OLAX*^^^^ 


IN  MEMORIAM 
Edmund  O'Neill 


A  TEXT-BOOK  OF 

EXPERIMENTAL  CHEMISTRY 

(WITH  DESCRIPTIVE  NOTES) 

LEE 


A  TEXT-BOOK  OF 

EXPERIMENTAL  CHEMISTRY 

(WITH  DESCRIPTIVE  NOTES) 

FOR  STUDENTS  OF 
GENERAL  INORGANIC  CHEMISTRY 


BY 

EDWIN    LEE 

PROFESSOR  OF  CHEMISTRY  IN  ALLEGHENY  COLLEGE 


WITH  57  ILLUSTRATIONS 


PHILADELPHIA 
P.  BLAKISTON'S  SON  &  CO, 

1012  WALNUT  STREET 
1908 


COPYRIGHT,  1908,  BY  P.  BLAKISTON'S  SON  &  CO. 
IN  MEMORIAM 


Printed  by 

The  Maple  Press 

York  Pa. 


PREFACE. 


WHILE  no  particular  claim  to  originality  is  made  for  this  text-book,  as 
many  of  the  experiments  have  been  described  previously,  yet  the  writer 
believes  that  the  book  will  be  found  to  be  something  more  than  a  mere 
compilation.  It  grew  originally  out  of  a  personal  demand  for  a  text- 
book which  would  embody:  (a)  a  clear,  accurate  and  comprehensive 
presentation  of  the  fundamentals  of  the  science;  (b)  specific  directions 
for  laboratory  work,  coupled  with  such  questions  as  lead  the  student  to 
observe,  compare  and  generalize,  and  wrould  therefore  provide  a  method 
for  the  scientific  development  of  the  principles  under  discussion;  (c)  a 
sufficient  amount  of  discussion  and  application  of  the  principles  involved 
in  the  experiments  to  foster  the  interest  and  to  direct  the  observa- 
tions that  energy  may  not  be  spent  indiscriminately,  and  (d)  those  physico- 
chemical  generalizations  which  are  essential  to  the  explanation  of  much 
of  the  phenomena  of  inorganic  chemistry. 

This  book  represents  an  endeavor  to  meet  these  requirements.  It 
is  not  intended  that  it  shall  take  the  place  of  a  large  descriptive  work  or 
the  instruction  of  the  teacher;  on  the  contrary,  it  is  designed  to  provide, 
primarily,  an  experimental  course  in  general  chemistry,  and  by  the  use  of 
"descriptive  notes"  and  questions  vitally  relate  it  to  the  lecture-room 
work.  It  is  scarcely  necessary  to  emphazise  the  importance  of  laboratory 
work  as  being  essential  to  a  thorough  comprehension  of  the  subject; 
but  this  same  work  has  a  very  doubtful  value  unless  it  is  carefully  directed 
and  correlated  with  the  lecture  and  text-book.  It  must  not  be  merely  a 
mechanical  part  of  the  course.  The  student  must  see  that  his  laboratory 
work  is  but  a  means  to  an  end — that  lectures  and  experiments  are  mu- 
tually helpful.  Very  frequently  the  laboratory  work  is  taught  too  much 
apart  from  the  course.  Beginners  often  complain,  and  more  frequently 
conduct  their  work  as  if  no  relation  existed  between  lectures  and  labora- 
tory work.  It  is  contended,  therefore,  that  the  laboratory  manual  should 
provide  something  to  make  obvious  this  relationship  and  to  assist  in  the 
fusion  of  the  two.  It  is  hoped  that  the  " notes"  appended  to  many  of  the 
experiments,  and  which  embody  discussions,  applications,  formulae,  etc., 
will  form  the  connecting  link  between  the  work  of  the  laboratory  and 


VI  PREFACE. 

the  work  of  the  lecture-room.  In  addition  to  the  foregoing,  sufficient 
amounts  of  descriptive,  theoretical  and  physical  chemistry  have  been 
incorporated  to  present  in  outline,  at  least,  the  essentials  of  general 
chemistry.  It  is  presumed,  however,  that  the  laboratory  work  will  be 
accompanied  by  a  full  course  of  lectures. 

The  classification  is  according  to  the  natural  families  of  the  periodic 
system. 

The  fundamental  concepts  of  the  science  are  first  built  up,  explained  and 
illustrated,  and  then  applied  persistently. 

Neither  the  inductive  nor  the  deductive  method  has  been  adhered 
to  with  "faddish"  tenacity.  Experience  has  taught  the  author  that  a 
combination  of  the  two  methods  can  be  relied  upon  to  yield  satisfactory 
results.  The  inductive  method  possesses  undoubted  virtues,  yet  experi- 
enced teachers  cannot  fail  to  appreciate  the  pertinency  and  significance 
of  Faraday's  remark,  "What  are  we  to  look  for,  Tyndall,"  as  the 
latter  was  about  to  perform  an  experiment.  A  too  vigorous  application 
of  even  this  method  usually  results  in  a  waste  of  energy.  There  is 
danger  in  any  extreme.  The  notes  and  questions  in  the  book  are  arranged 
with  a  view  to  imparting  definiteness  of  purpose  and  aim  to  the  labora- 
tory work. 

Most  of  the  exercises  have  been  used  in  the  type-written  form  for 
a  number  of  years.  They  cannot  be  regarded  as  "impracticable  and 
beyond  the  capabilities  of  the  average  student."  An  effort  has  been 
made  to  lighten  the  labor  of  the  teacher  by  making  the  directions  specific. 
Further,  a  series  of  graduated  questions  have  been  incorporated  among 
these  directions  in  order  to  asisst  the  student  in  correlating  and  generaliz- 
ing. This  also  should  save  the  teacher's  time  and  energy.  The  rigid- 
ness  imparted  by  the  use  of  specific  directions  is  overcome  by  the  intro- 
duction of  abundant  material  to  meet  varied  conditions — a  much  larger 
number  of  experiments  being  inserted  than  will  likely  be  used  either  in 
lecture-room  or  laboratory.  This  enables  the  teacher  to  make  a  selection. 

No  hesitancy  has  been  manifested  in  introducing  physico-chemical 
generalizations  wherever  they  have  seemed  necessary  to  rational  correla- 
tion and  explanation  of  facts.  The  importance  of  these  generalizations 
in  connection  with  the  teaching  of  general  chemistry  is  now  generally 
recognized,  yet  the  writer  confesses  that  he  knows  of  no  text-book  on 
"experimental  chemistry"  which  has  been  brought  abreast  of  the  times 
by  appropriating  and  incorporating  the  results  of  the  labors  of  the  physical 
chemist.  It  may  be  true  that  physical  chemistry  cannot  and  should  not  be 


PREFACE.  VI] 

taught  in  the  first  year  of  chemistry,  but  in  the  opinion  of  the  author  and 
at  least  one  other,  "many  of  these  generalizations  which  have  been  brought 
in  by  this  new  physical  chemistry,  and  which  affect  fundamentally  the 
whole  science  of  chemistry,  can  and  should  be  introduced  into  general 
chemistry."  The  author  wrould  not  leave  the  impression  that  the  older 
generalizations  and  methods  of  developing  the  subject  have  been  aban- 
doned; on  the  contrary,  they  have  been  retained  and  an  effort  made  tc 
rationalize  them  by  supplementing  them  with  more  recent  data  taken 
from  the  domain  of  physical  chemistry.  The  book  represents,  therefore, 
an  effort  to  fuse  modern  views  and  recent  advances  with  those  older  views 
which  have  stood  the  test,  and  to  explain  chemical  phenomena  in  a  man- 
ner that  is  in  accord  with  modern  chemical  thought. 

The  role  of  energy  in  chemical  reaction  has  been  given  unusual  promi- 
nence. The  elements  of  thermochemistry  have  been  presented.  More 
space  has  been  allotted  to  "solutions,"  "the  ion  theory,"  "chemical 
equilibrium"  and  the  "mass  law"  than  is  usually  given  in  books  of  this 
character.  It  is  believed,  however,  that  the  results  obtained  warrant 
this.  A  brief  comprehensive  statement  of  the  historical  development 
of  "the  electrolytic  dissociation  theory"  has  been  given  that  the  student 
may  become  familiar  with  the  story  of  the  gradual  development  of  at 
least  one  chemical  theory.  The  author  has  observed  that  students  take  a 
keen  delight  in  stating  just  how  much  was  contributed  to  the  develop- 
ment of  a  chemical  theory  by  this  or  that  chemist.  With  a  similar  purpose 
in  view,  the  writer  has  quoted  references  verbatim,  rather  than  record  the 
gist  of  them. 

Chapter  VII  presents  in  outline  such  subjects  as  the  kinetic-molecular 
hypothesis,  atomic  theory,  valence,  formulae,  equations  and  units.  This 
gives  flexibility  to  the  course  by  allowing  the  teacher  to  develop  the 
several  subjects  in  a  manner  best  suited  to  the  methods  in  use. 

The  consulting  of  easily-found  references  is  demanded  frequently. 
The  London  Chemical  News  says  in  this  connection:  "The  habit  of 
looking  up  whatever  needs  greater  elucidation  augurs  well  for  good 
culture  and  education."  If  possible,  many  of  the  books  mentioned  in 
the  reference  list  should  be  made  accessible  to  the  students. 

The  author's  supreme  purpose  has  been  to  present  a  system  of  chemistry 
rather  than  an  assemblage  of  chemical  facts,  and  to  build  a  book  which 
will  aid  the  student  in  cultivating  correct  habits  of  thought. 

The  author  acknowledges  with  pleasure  the  help  and  inspiration  of 
those  masterful  teachers  with  whom  he  has  been  permitted  to  work. 


Vlll  PREFACE. 

To  his  fellow  teachers  who  have  been  generous  with  valuable  sugges- 
tions, he  acknowledges  his  indebtedness.  The  writer  desires  especially 
to  thank  Professor  Arthur  B.  Lamb,  of  New  York  University,  and 
Dr.  J.  H.  Mathews  of  the  University  of  Wisconsin  for  helpful  criticisms. 

For  courtesies  received,  the  writer  hereby  extends  his  thanks  to 
Professor  T.  W.  Richards  of  Harvard  University,  Professor  A.  A.  Noyes 
and  Dr.  G.  N.  Lewis  of  Massachusetts  Institute  of  Technology,  Professor 
Harry  C.  Jones  and  Dr.  H.  N.  Morse  of  Johns  Hopkins  University,  Professor 
T.  L.  R.  Morgan  of  Columbia  University,  Professor  Alexander  Smith  of 
the  University  of  Chicago,  and  Professor  F.  A.  Gooch  of  Yale  University. 

I  am  indebted  to  my  publishers,  P.  Blakiston's  Son  &  Company, 
for  the  use  of  certain  illustrations  from  several  of  their  publications 
and  desire  to  express  my  appreciation  of  their  courtesy. 

The  author  will  be  grateful  to  those  who  report  errors  and  ambiguity. 

ALLEGHENY  COLLEGE,  EDWIN  LEE. 

Meadville,  Penna.,  June,  1908. 


CONTENTS. 


INTRODUCTORY. 

PAGE 

CHAPTER  I. 
FUNDAMENTAL  CONCEPTS 17 

Relationships  and  Definitions. — Entity  Defined. — Physics  versus 
Chemistry. — Note  on  Matter. — "System"  Defined. — Specimens  or 
Bodies. — Substances. — Homogeneity  and  Heterogeneity. — Proper- 
ties.— Identity. — Experiments. — Note  on  Energy. — Forms  of  Energy. 
— First  Law  of  Energy. — Transformation  of  Energy. — Experiments . — 
Relation  of  Energy  to  Properties  of  Matter. — Phase. — Relation  of 
Energy  to  Phase  of  a  Substance. — Experiments. 

CHAPTER  II. 

CHARACTERISTICS  OF  PHYSICAL  AND  CHEMICAL  CHANGES 26 

Note. — Definitions. — Physical  and  Chemical  Changes. — Electrolysis 
of  Water. — Change  of  Identity. — Analysis. — Effect  of  Heat  on  Mer- 
curic Oxide. — Effect  of  Heating  Copper  in  Air. — Synthesis. — Solution 
Facilitates  Chemical  Action. — Note  on  Surface  Phenomena. — Relation 
of  Energy  to  Chemical  Change. — Exercises. 

CHAPTER  III. 

ELEMENTARY  SUBSTANCES,  MIXTURES  AND  COMPOUNDS 31 

Note. — Second  Law  of  Matter. — Examination  of  the  Elementary  Sub- 
stances.— Energy  Content  Affects  Properties  of  Elementary  Sub- 
stances.— Mixtures — Heterogeneous  and  Homogeneous. — Note. — 
Mixtures  and  Compounds. — Definitions. — Interaction  of  Compounds. 
— Metathesis. — Reaction. — Reagents. — "Factors."— "Products." 

CHAPTER  IV. 

NOTE  ON  THE  ENERGETICS  OF  CHEMISTRY 34 

Theory. — First  Law  of  Energy. — "Matter  Content." — "Energy 
Content." — Internal  Energy,  Cohesion  Energy,  Disgregation  Energy, 
Chemical  Energy. — Free  and  Bound  Energy. — LeChatelier's  Theorem. 
— Application  of  Theory. — Red  and  Yellow  Phosphorus. — Allotropy. — 
Thermochemistry. — Endothermic  and  Exothermic  Reactions. — First 
Law  of  Thermochemistry. — Law  of  Hess. — Heat  of  Reaction  versus 
Heat  of  Formation. 

CHAPTER  V. 

SCIENCE,  ITS  METHODS  OF  DEVELOPMENT, — CLASSIFICATION 40 

Definitions  of  Phenomenon,  Experiment,  Fact,  Law,  Rule. — Induc- 
tion, Deduction. — Hypothesis,  Theory,  Prediction. — Scientific  versus 
Systematic  Methods. — Physical  Sciences  and  Biological  Sciences. — 
Abstract  and  Concrete  Sciences. 


x  CONTENTS. 

CHAPTER  VI. 

PAGE 

QUANTITATIVE  RELATIONSHIPS.     LAWS  AND  CHEMICAL  EQUIVALENTS  ...  43 

Note- — Conservation  of  Mass  in  a  Limited  System. — Analysis  of 
Water. — First  Law  of  Matter. — Combination  in  Definite  Propor- 
tions by  Weight. — Synthesis  of  Copper  Oxide,  Magnesium  Oxide, 
Iron  Oxide. — Law  of  Definite  Proportions. — Constancy  of  Composi- 
tion.— Combination  of  Two  Substances  in  Varying  Proportions  by 
Weight. — Law  of  Multiple  Proportions. — Combining  Volumes. — 
Synthesis  of  Water  from  Hydrogen  and  Oxygen  Gases. — Syntheses  of 
Steam. — The  Law  of  Gay-Lussac. — Law  of  Reciprocal  Proportions. — 
Chemical  Equivalents. — Equivalent  Weight  of  Zinc  by  Displacement 
of  Hydrogen. 

CHAPTER  VII. 

OUTLINES,  THEORIES,  FORMULAE,  VALENCE,  EQUATIONS,  UNITS 50 

Structure  of  Matter. — Molecule. — Kinetic-Molecular  Hypothesis,  and 
Its  Application. — Critical  Phenomena. — Combining  Volumes. — 
Atom. — Atomic  Theory  and  Its  Application. — Electronic  Theory  of 
Matter. — Nomenclature  of  Atomic  Theory. — Simple  and  Compound 
Molecules. — Molecular  Weight. — Mass  and  Energy  Equations. — 
Note. — Valence. — Multiple  Valence. — Kinds  of  Formulae. — Units  of 
Measurement. — Stoichiometry. — Methods  for  the  Determination  of 
Molecular  and  Atomic  Weights. — References. — Problems. 

CHAPTER  VIII. 

OXYGEN 54 

Introductory  Note  on  Occurrence  and  Properties  of  Oxygen. — Sources 
of  Oxygen. — Laboratory  Sources  of  Oxygen. — Oxidation  of  Carbon, 
Phosphorus,  Sulphur,  Iron. — Oxides. — "Ous"  and  "ic"  Endings. — 
Kindling  Temperature. — Note  on  Heat  of  Combustion. — Energy 
Equations. — Catalytic  Action  of  Manganese  Dioxide. — Note  on 
Catalyzers. — Preparation  and  Properties  of  Ozone. — Note  on  Energy 
of  Oxygen. — Oxygen  and  Ozone. — Determination  of  the  Density 
Relations  of  Determination  of  the  Volume  of  Oxygen  Liberated  from 
a  Given  Weight  of  Potassium  Chlorate. — Percentage  of  Oxygen  in 
the  Air. — Dalton's  Law  of  Partial  Pressures. — Problems. 


CHAPTER  IX. 

HYDROGEN 64 

Note. — Methods  for  Preparing  Hydrogen. — Displacement  of  Hydro- 
gen from  Acids  by  Metals. — Note  on  Change  of  Energy  in  a  System. — 
Properties  of  Hydrogen. — Effusion  of  Hydrogen. — Graham's  Law  of 
Effusion. — Occlusion  of  Hydrogen. — Catalytic  Action  of  Spongy 
Platinum. — Note  on  Catalyzers  and  Surface  Phenomena. — Reduction 
of  Cupric  Oxide. — The  Nascent  State. — Nascent  Hydrogen. — Prepara- 
tion and  Properties  of  Hydrogen  Dioxide. — Hydrogen  Equivalent  of 
Zinc. — Outline  of  Hydrogen. — TJaeHttochernical  Deportment  of  Hydro- 
gen and  Oxygen  in  Formation  of  Water  and  Hydrogen  Dioxide. — 
Problems. 


CONTENTS.  XI 

CHAPTER  X. 

PAGE 

WATER 75 

Note. — Composition  and  Properties  of  Water. — Purity. — Temper- 
ature.— Density  Graph  of  Water. — Hydrolysis. — Chemical  Union  with 
Oxides. — Water  of  Hydration. — Vapor  Tension  of  Substances,  Efflor- 
escence and  Deliquescence. — Determination  of  Water  of  Crystalliza- 
tion.— Outline  of  Properties  of  Water. — Note  on  Polymerization  and 
Hydrolysis. 

CHAPTER  XI. 

SOLUTIONS 82 

Note  and  Definitions. — Kinds  of  Solutions. — Solubility  of  Solids. — 
Characteristics  of  a  Solution. — Note  on  Colloidal  Solutions. — Surface 
and  Diffusion  Phenomena. — Solution  Tension. — Note  on  Equilibria  in 
Solutions. — Osmotic  Pressure. — Pfeffer's  Data. — Morse  and  Frazer's 
Results. — Van't  Hoff's  Law. — Effect  of  Temperature  on  Solubility  of 
Solids. — Supersaturated  Solutions. — Thermal  Phenomena. — Relative 
Solubility  of  Solids. — Terminology  of  Solutions. — Table  of  Solubility 
of  Solids. — Note  on  Heat  of  Solution. — Table,  Heat  of  Solution  of 
Solids,  Liquids  and  Gases. — Effect  of  Pressure  on  Solubility. — Solu- 
bility of  Pairs  of  Liquids. — Relative  and  Mutual  Solubility. — Effect  of 
Temperature. — Critical  Solution  Temperature. — Distribution  of 
Solute  Between  Two  Immiscible  Solvents. — Law  of  Distribution. — 
Solubility  of  Gases. — Pressure  of  Gas  Affects  Solubility. — Henry's 
Law. — Effect  of  Solvent  on  Solubility. — Effect  of  Temperature  on 
Solubility  of  Gases. — Elevation  of  the  Boiling  Point. — Depression  of 
the  Freezing  Point. — Notes. — Raoult's  Laws. — Lowering  of  the  Solu- 
tion Tension. — Note. — Problems. 

CHAPTER  XII. 

ACIDS,  BASES  AND  SALTS 104 

Chlorine. — Preparation  and  Experimental  Study  of  Its  Properties. — 
Note. — Hydrogen  Chloride  Gas. — Acids. — General  and  Specific 
Properties  of  Acids. — Anhydrides. — Acid  Radicals. — Classification 
of  Acids:  Sodium. — Its  Properties. — Note. — Sodium  Hydroxide. — 
Bases — General  and  Specific  Properties. — Classification  of  Bases. — 
Interaction  of  Acids  and  Bases. — Neutralization.: — Salts. — Classifica- 
tion and  Nomenclature  of  Salts.— Metal  and  Non-Metal  Defined. — 
Problems. 

CHAPTER  XIII. 

CHEMICAL  EQUILIBRIUM,  REVERSIBLE  REACTIONS,  MASS  LAW,  DISSOCIATION,  1 19 

Note  on  Chemical  Equilibrium  and  Law  of  Concentration  Effect. — 


Mathematical   Form   of   Mass   Law. — Affinity   Constant. — Reactions 

are   Incom- 
a  Gas. 


Which  are  Approximately  Complete. — Reactions   which   ; 
plete. — Equilibrium. — Berthollet's  Laws. — Dissociation  of 


CHAPTER  XIV. 

NOTE  ON  THE  MODERN  THEORY  OF  SOLUTION 124 

Historical  Development  of  the  Electrolytic  Dissociation  Theory. — 
Statement  of  the  Theory. — Le  Blanc,  Nernst,  Thomson  and  Beilby, 
quoted  in  regard  to  the  probable  source  of  the  electric  charges  on  the 
ions. — Application  of  Theory. — Electrolysis. — Electro-Chemical  Equiv- 
alent.— Names  of  Ions. — Electrolytes,  Half-Electrolytes  and  Non- 


Xll  CONTENTS. 

PAGE 

Electrolytes. — Note. — lonization  in  Solution. — Effect  of  the  Nature  of 
the  Solvent  and  the  Solute  on  lonization. — Effect  of  Dilution. — 
Ostwald's  Law  of  Dilution. — Dissociation  Constant. — Color  of  Ions. — 
Chemical  Conduct  of  Ions. — "Strength"  of  Acids  and  Bases. — 
Tables. — Thermochemical  Support  of  the  Dissociation  Theory. — 
Change  in  Volume  Support. — Salts. — Compound  Salts  versus  Complex 
Salts. — Hydrolysis. — Ionic  Equilibrium. — "Ion  Product." — "Solu- 
bility Product." — Solution  Tension. — Relation  of  Chemical  Energy 
and  Electrical  Energy. — Note. — E.M.F.  Is  a  Measure  of  Chemical 
Energy  When  the  Process  is  Reversible. — Table  of  Single  Potentials. — 
Heat  of  lonization. — Table. 

THE  NON-METALS  OR  ACID-FORMING  ELEMENTS. 
CHAPTER  XV. 

THE  CHLORINE  FAMILY 152 

Introductory  Note. — Preparation  and  Properties  of  Fluorine. — Hydro- 
gen Fluoride. — Silicon  Tetrafluoride. — Note:  Chlorine,  Preparation 
and  Properties. — Hydrogen  Chloride. — Chlorides. — Oxygen  Deriva- 
tives of  Chlorine. — Chlorates  and  Perchlorates :  Bromine,  Prepara- 
tion and  Properties. — Substituting  Power  of  Chlorine  as  Compared 
with  Bromine. — Hydrogen  Derivatives. — Bromides. — Oxygen  Deriva- 
tives.— Bromates,  etc.:  Iodine,  Preparation  and  Properties. — Tinc- 
tures.— Substituting  Power  of  Chlorine  and  Bromine  as  Compared  with 
Iodine. — Liberation  of  Iodine  by  Oxidizing  Agents. — Displacement 
of  Sulphur  from  Hydrogen  Sulphide  by  Chlorine,  Bromine  or  Iodine. — 
Hydrogen  Derivatives  of  Iodine. — Iodides. — Oxygen  Derivatives. — 
lodates.  etc.:  Periods  of  Induction  in  Chemical  Reactions. — Resem- 
blance of  the  Properties  of  Corresponding  Compounds  of  the  Halogens. 
— Detection  of  Fluorides,  Chlorides,  Bromides,  and  Iodides,  together 
in  a  Solution. — Comparative  Table. — Note. 

CHAPTER  XVI. 

CLASSIFICATION  OF  THE  ELEMENTS 167 

Note. — Prout's  Hypothesis. — Dobereiner's  Triads. — Newland's  Law 
of  Octaves. — Mendeleef's  Original  Table. — Lothar  Meyer's  Natural 
Classification. — Periodic  Arrangement. — Bayley's  Table. — Atomic 
Volume  Table. — Table  of  Atomic  Volume  and  Specific  Gravity. — 
Exercises. 

CHAPTER  XVII. 

GROUP  VI.— FAMILY  M.— THE  OXYGEN  FAMILY 176 

Note. — Oxygen. — Hydrogen  Derivatives  and  Their  Thermal  Con- 
duct.— Allotropy. — Ozone :  Sulphur. — Allotropy. — Note. — Sulphides. 
— A  Test  for  Sulphides. — Oxygen  Derivatives  of  Sulphur. — Sul- 
phuric Acid. — Manufacture  of  Sulphuric  Acid. — Soluble  and 
Insoluble  Sulphates. — Contraction  in  Volume  When  Sulphuric  Acid 
and  Water  are  Mixed.— Table:  Selenium:  Tellurium:  Compara- 
tive Table. — Problems. 

CHAPTER  XVIII. 

NITROGEN  AND  THE  ATMOSPHERE. — THE  HELIUM  FAMILY 192 

Nitrogen. — Preparation  and  Properties. — Determination  of  the  Con- 
stituents of  the  Atmosphere.— Weight  of  a  Liter  of  Air:  The  Helium 
Family. — Note. — Tables. — Problems. 


CONTENTS.  Xlll 

CHAPTER  XIX. 

PAGE 

GROUP  V. — FAMILY  M. — NITROGEN  FAMILY 199 

Note  on  Relationships. — Nitrogen,  Preparation  and  Properties. — 
Hydrogen  Derivatives. — Ammonia. — Ammonium  Salts. — Nitrogen  and 
the  Halogens. — Preparation  of  an  Endothermic  Compound,  Nitrogen 
Tri-iodide. — Composition  of  Ammonia  Gas. — Oxides  of  Nitrogen  are 
Endothermic  Compounds. — Note  on  Heat  of  Formation. — Nitric 
Acid. — Note  and  Equations  Relative  to  the  Conduct  of  Nitric  Acid. — 
Aqua  Regia. — Nitrates:  Phosphorus. — Allotropy. — Hydrogen  Deri- 
vatives.— Phosphorus  and  the  Halogens. — Oxygen  Derivatives. — 
Phosphoric  Acid. — Tests  to  Distinguish  the  Phosphates:  Arsenic. — 
Properties. — Derivatives. — Tests  for  Arsenic:  Antimony. — Deriva- 
tives.— Hydrolysis  of  Antimony  Trichloride. — Sulphides  of  Antimony: 
Bismuth. — Properties. — Alloys. — Hydrolysis  of  Bismuth  Nitrate. — 
Comparative  Table. — Problems. 

CHAPTER  XX. 

CARBON  FAMILY 226 

Primary  Group. — Carbon. — Allotropy. — Carbon  and  Oxygen. — 
Energy  Equations. — Carbonates. — Application  of  the  Law  of  Hess. — 
Carbon  and  Hydrogen  (Hydrocarbons). — Acetylene. — Coal  Gas. — 
Carbon  Disulphide. — Carbon  and  the  Halogens. — Carbon  and  Nitro- 
gen.— Cyanides. — A  study  of  Flames. — Temperature  of  Flame  of 
Bunsen  Burner:  Silicon. — Note. — Silic  Acid. — Dialysis. — Fluo- 
silicic  Acid. — A  Test  for  Silicates. — Note. — Comparative  Table. — 
Problems. 

CHAPTER  XXI. 

SOME  COMMON  CARBON  COMPOUNDS 246 

Introductory  Note. — Hydrocarbons. — Alcohols. — Phenols. — Ethers. — 
Aldehydes. — Acids. — Ketones. — Halides  of  Carbon. — Esters. — Ammo- 
nia.— Derivatives. — Organo-Mineral  Compounds. — Carbohydrates. — 
Glucosides. — Albuminoids. — Isomerism. — Flashing-point  of  Kerosene. 
—Tests  for  Ethyl  Alcohol. — Glycerin. — Soaps. — Carbolic  Acid. — 
Glacial  Acetic  Acid. — Tartaric  Acid. — Oxalic  Acid. — Chloroform. — 
lodoform. — Butter. — Fatty  or  Fixed  Oils. — Essential  or  Volatile  Oils 
— Sugars. — Starches. — Gun-Cotton  or  Smokeless  Powder. — Paper. 

CHAPTER  XXII. 

ALUMINUM  FAMILY       • 271 

Note. — Boron. — Properties. — Borax. — Solubility  of  Metallic  Oxides  in 
Borax.— Borax.— Bead  Tests.— Table.— Detection  of  Boron.— Com- 
parative Table. 

THE  METALS  OR  BASE-FORMING  ELEMENTS. 
CHAPTER  XXIII. 

INTRODUCTORY  NOTE 276 

Physical  Properties  of  the  Metals. — Chemical  Properties  of  the  Metals. 
— Occurrence  of  the  Metals  in  Nature. — Table. — Classification  of  the 
Metals. 


XIV  CONTENTS. 

CHAPTER  XXIV. 

PAGE 

GROUP  I. — FAMILY  M.     ALKALI  METALS 280 

Introductory  Note. — Lithium. — Flame  Color  of  Lithium  Compounds : 
Sodium. — Experiments. — Tests. — Alkalimetry  and  Acidimetry:  Po- 
tassium.— Experiments. — Tests. — Gunpowder:  Ammonium. — Experi- 
ments.— Detection  of  the  Alkali  Metals  in  a  Mixture:  Rubidium: 
Caesium:  Problems. 

CHAPTER  XXV. 

GROUP  I. — FAMILY  M.     THE  COPPER  FAMILY 292 

Introductory  Note. — Copper. — Experiments. — Scheme  of  lonization 
for  Double  Salts  and  Complex  Salts. — Note:  Silver. — Experiments. — 
Photography:  Gold. — Experiments. — "Purple  of  Cassius:"  Prob- 
lems. 

CHAPTER  XXVI. 

GROUP  II. — FAMILY  M.     ALKALINE  EARTH  METALS 306 

Introductory     Note. — Relative     Solubility     of     Salts. — Glucinum: 
Magnesium. — Experiments :     Calcium. — Experiments. — Quicklime. — 
Plaster  of  Paris. — "Permanent  Hardness  of  Water:"     Strontium. — 
Experiments:     Barium. — Experiments:     Detection    of    the    Alkaline 
Earth  Metals  in  a  Mixture:     Radium:     Exercises. — Problems. 

CHAPTER  XXVII. 
SECONDARY  FAMILY  OF  THE  ALKALINE  EARTHS 316 

Introductory  Note. — Zinc. — Comparative  Table. — Experiments, — 
An  Experimental  Study  of  Ionic  Equilibrium  and  "Concentration 
Effect:"  Cadmium. — Formation  of  Salts  and  Analytical  Reactions: 
Mercury. — Properties. — Formation  of  Salts  and  Analytical  Reactions: 
Problems. 

CHAPTER  XXVIII. 

THE  ELEMENTS  OF  GROUP  III 326 

Family   M — Scandium,   Yttrium,    Lanthanum,   Ytterbium. — Note. — 

Absorption  and  Emission  Spectra. — Eka-boron  and  Eka-Aluminum  of 

Mendeleeff. 

Family  m — Note. — Boron. — Aluminum. — Experiments. — "Thermit." 

— Ruby. — Sapphire. — Emery. — Clay. — Classification       of        Alums: 

Thallium. — Properties. — Salts  of  Thallium:     Problems, 

CHAPTER  XXIX. 

THE  ELEMENTS  OF  GROUP  IV 333 

Family  M — Titanium,  Zirconium,  Cerium,  Thorium. — Note. 
Family  m — Note. — Tin-plate.  — Type-metal. — Shot-metal. — German- 
ium: Tin. — Experiments. — Determination  of  the  Equivalent  Weight 
of  Tin. — Analytical  Reactions:  Lead. — "Lead  Tree." — Formation 
and  Preparation  of  Salts. — Chemical  Principles  Involved  in  the  Manu- 
facture of  White  Lead :  Problems. 

CHAPTER  XXX. 

ELEMENTS  OF  GROUP  V 343 

Family  M —  Vanadium,  Columbium,  Tantalum. — Note. 


CONTENTS.  XV 

CHAPTER  XXXI. 

PAGE 
ELEMENTS  or  GROUP  VI 344 

Family  M — Chromium. — Note. — Experiments. — Chromyl  Chloride, 
a  Test  for  Chlorides. — Analytical  Reactions:  Molybdenum. — Ammo- 
nium Phosphomolybdate :  Tungsten. — Tungsten  Steel:  Uranium. — • 
Pitchblende. — Radio-activity :  Problems. 

CHAPTER  XXXII. 

ELEMENTS  OF  GROUP  VII 351 

Family  M — Manganese. — Note. — Experiments. — Equations. — Analyt- 
ical Reactions:  Problems. 

CHAPTER  XXXIII. 

TRANSITION  ELEMENTS 355 

First  Long  Period. — Note:  Irori. — Note. — Application  of  Iron  to 
Industrial  Arts. — "Passive  State." — Corrosion  of  Iron. — Experiments. 
Reactions  of  Ferrous  and  Ferric  Salts. — Ferrocyanides  and  Ferricy- 
anides. — Oxidation  and  Reduction  of  Iron  Salts:  Cobalt. — Note. — 
Experiments. — Reactions :  Nickel. — Note. — Electro-plating. — Alloys : 
Problems. 

CHAPTER  XXXIV. 

TRANSITION  ELEMENTS 366 

Second    Long    Period. — Note:     Ruthenium. — Note:     Rhodium. — 
Note:     Palladium. — Note. — Occlusion — "Solid     Solutions." — Fourth 
Long  Period. — Note:    Osmium. — Note:     Iridium. — Note:    Platinum. 
— Note. — Platinum  Black. — Experiments. — Platinum  as  a  Catalyzing 
Agent. — Reactions. 

CHAPTER  XXXV. 

RELATIONS  WITHIN  THE  GROUPS  OF  THE  PERIODIC  CLASSIFICATION  ....          370 
Grouping  of  the  Metals  (Cations)  for  Purposes  of  Analysis;  Grouping 
of  Non-Metallic  Radicals  (Anions)  for  Purposes  of  Analysis. 

APPENDIX  I 373 

PRELIMINARY  EXERCISES 

The  Bunsen  Burner. — The  Blast-Lamp. — Manipulation  of  Glass. — 
Perforation  of  Stoppers. — Treatment  of  Rubber  Corks  and  Tubing. — 
The  Wash  Bottle. — Measuring  Instruments. — Calibration  by  Weigh- 
ing.— Problems 

APPENDIX   II  (Tables) 381 

REFERENCE  BOOKS 415 

LABORATORY  INSTRUCTIONS xx 

INDEX    ...                                                                             419 


LIST  OF  TABLES. 


PAGE. 

Osmotic  Pressure  of  Cane  Sugar 86 

Osmotic  Pressure  and  Molecular  Weights 87 

Solubility  of  Solids 91 

Heat  of  Solution  of  Solids 92 

Heat  of  Solution  of  Liquids 92 

Heat  of  Solution  of  Gases 93 

Solubility  of  Gases 96-97 

Constants  (Elevation  of  the  Boiling  Point) 99 

Raoult's  Tables  (Depression  of  the  Freezing  Point) 101 

Names  of  Ions 135 

Velocity  of  Ions 135 

Heat  of  Neutralization 143 

Acids,  Bases,  Salts. — Per  Cent.  Ionized 145 

Order  of  Solution  Tension  of  the  Metals 147 

Single  Potentials 150 

Heat  of  lonization 151 

Comparative  Table.     The  Halogens 165 

List  of  Non-Metals 167 

Triads   of   Dobereiner 168 

Newland's     Octaves 168 

Mendeleef's  Original  Table 169 

Lothar  Meyer's  Original  Table 170 

Periodic  arrangement: 

Table     1 171 

Table  II 172 

Bayley's    Table 173 

Atomic  Volume,  Specific  Heat,  Valence 1 74 

Atomic  Volume — Specific  Gravity 175 

Comparative  Table.     The  Oxygen  Family 191 

Comparative  Table.     The  Helium  Family 197 

Comparative  Table.     The  Nitrogen  Family 225 

Temperature  of  Bunsen  Burner  Flame 240 

Comparative  Table.     Carbon  and  Silicon 244 

Colors  of  Borax  Beads 274 

Comparative  Table.     Boron  and  Aluminum 275 

Comparative  Table.     Zinc,  Cadmium  and  Mercury 317 

Comparative  Table.     Germanium,  Tin  and  Lead 337 

xvii 


XViii  LIST    OF    TABLES. 

PAGE 

Grouping  of  the  Metals  (Cations)  for  Purposes  of  Analysis • 370 

Grouping  of  the  Non-Metallic  Radicals  (Anions)  for  Purposes  of  Analysis .    .    .    .371 

Metric  Measures  with  English  Equivalents 381 

English  Measures  with  Metric  Equivalents 382 

Conversion  of  Thermometric  Readings 383 

Corrections  for  Mercury-in-Glass  Thermometers 383 

Density  of  Water 384 

Volume  of  One  Gram  of  Water  at  Various  Temperatures 385 

Density  of  Mercury 385 

Barometric    Corrections 386 

Reduction  to  Vacuum  of  Weighings  Made  in  Air 387 

Correction  Factors  for  Calibrating  Glass  Vessels 388 

Vapor  Tension  of  Water  (o°-35°) 389 

Vapor  Tension  of  Water  (o°-ioo°) 390 

Vapor  Pressure  of  Mercury 391 

Boiling  Temperature  (t)  of  Water  at  Barometric  Pressure  (b) 392 

Boiling  Temperature  (t)  of  Water  at  Pressure  of  (a)  Atmospheres 392 

Physical  Constants 393 

Composition  of  the  Air  by  Volume 394 

Diffusion  of  Gases 394 

Specific  Heat — Atomic  Heat 395 

Heat  of  Formation 

(a)  Chlorides,   Bromides,   Iodides 397 

(b)  Sulphates,  Nitrates,  Carbonates 398 

(c)  Oxides,  Hydroxides,  Sulphides 399 

Solubilities 400 

Solubilities  of  Salts  (Quant.) 401 

Composition  of  Some  of  the  Important  Alloys 402 

Scale  of  Hardness 404 

Color  Scale  of  Temperature 404 

Indicators 405 

Soap-Bubble  Solutions 405 

Graduated  Solutions: 

(a)  Fehling's  Solution 406 

(b)  Nessler's  Solution 406 

Percentage  and  Specific  Gravity  of  Solutions  at  15°  C.: 

(a)  Sulphuric     Acid 407 

(b)  Hydrochloric  Acid 408 

(c)  Nitric   Acid 408 

(d)  Ammonium     Hydroxide 409 

Proportion  by  Weight  of  Absolute  Alcohol 410 

Proportion  by  Volume  of  Absolute  Alcohol 411 

List  of  Elements,  Valences,  Discoverer,  etc 412-414 


LABORATORY  INSTRUCTIONS. 


After  receiving  assignment  of  desk,  procure  key  from  rack.     Check  your  appara- 
tus by  comparison  with  list  found  in  drawer.     Report  any  differences  at  once  to  the 
assistant. 

Read  the  "Regulations"  posted  in  various  places  in  the  laboratory.  Always  read, 
in  entirety,  the  directions  for  performing  an  experiment  before  assembling  apparatus. 

It  is  well  to  provide  yourself  with  apron  or  blouse  to  protect  clothing  while  at  work , 
also  white  soft  cloth  to  be  used  for  wiping  apparatus.  A  sponge  is  convenient. 

Procure  a  note-book  (Instructions).     Make  a  neat,  permanent  and  true  report  of 
each  experiment  immediately  after  its  performance  under  the  following  heads: 
I.  Object  of  Experiment. 
II.  Manipulation.     (What  you  did.) 

III.  Observed  Phenomena. 

IV.  Conclusions  or  Results. 

V.  Give  equations  if  requested. 
VI.  Answer  questions. 

VII.  Errata — (any  mistakes  should  be  noted  under  this  head). 
VIII.  A  diagram  of  apparatus  frequently  facilitates  the  interpretation  of  an 
experiment. 

Students  must  work  independently,  both  as  to  manipulation  and  records,  unless 
otherwise  specified. 

When  weighing  is  necessary,  use  rough  balances  unless  experiment  is  marked 
"Quant." 

Do  not  carry  the  bottles  containing  the  various  substances  from  the  shelves  to  your 
desk.  Keep  the  bottles  in  order.  Use  a  test-tube  or  watch  glass  or  a  piece  of  paper 
to  transfer  substances.  Do  not  return  to  the  bottles  unused  portions  of  substances 
unless  you  have  secured  permission  from  the  instructor.  If  quantity  to  be  used  is 
not  specified  use  2  cm.3  or  3  cm.3 

Do  not  throw  anything  away  until  quite  sure  that  you  are  through  with  it.  Throw 
waste  liquids  into  sink;  other  waste  material  into  waste  jars. 

When  replacing  apparatus  use  order-sheet  (pink). 

T.  O.  means  that  apparatus  for  experiment  is  to  be  secured  temporarily  from  the 
instructor;  i.e.,  it  is  to  be  returned  to  him  after  performance  of  experiment. 

L.  T.  suggests  that  experiment  can  be  performed  to  advantage  on  lecture -table. 

Examine  your  desk  on  entering  the  laboratory.  If  anything  has  been  disturbed  or  is 
missing,  report  same  at  once  to  assistant. 

Before  leaving  the  laboratory,  place  your  desk  in  order.  Cleanliness  is  absolutely 
necessary.  Failure  to  observe  this  contributes  to  failure.  After  observing  these  instruc- 
tions, lock  your  desk  and  place  key  on  rack.  See  that  gas  and  water  are  turned  off. 

NOTE. — The  student  will  be  graded  on  his  "laboratory  deportment";  i.  e.,  the 
persistency  with  which  the  foregoing  instruct  ons  are  adhered  to. 


XX 


A  TEXT-BOOK 
OF 

EXPERIMENTAL  CHEMISTRY. 


CHAPTER  I. 
FUNDAMENTAL  CONCEPTS: 

RELATIONSHIPS    AND    DEFINITIONS. 

Through  the  medium  of  our  senses  we  are  constantly  receiving  sensa- 
tions which  we  interpret  objectively,  i.e.,  we  locate  the  cause  of  a  sensation 
in  a  particular  portion  of  space.  As  the  result  of  our  experience  we  assume 
that  the  physical  universe  has  an  objective  existence,  and  that  our  ac- 
quaintance with  it  depends  solely  upon;  otic  senses.  {  ' 

Men  have  given  the  name  "thing."  or  "entity*'  to  the'cause  of  a  sen- 
sation— to  that  which  has  the  objective-  exisfe'uojr:  '-^he' entities  with 
which  the  scientist  is  particularly  concerned  are  Matter  and  Energy. 
Time,  Temperature,  Space,  Velocity,  etc.,  are  not  things.* 

It  will  be  readily  recalled  from  the  student's  experience  in  the  study  of 
physics  that  the  two  classes  of  things  with  which  the  physicist  has  to  do 
are  those  previously  enumerated.  The  same  is  true  of  the  chemist. 

The  fundamental  difference  between  Physics  and  Chemistry  rests 
upon  the  relative  amount  of  emphasis  laid  upon  either  Energy  or  Matter. 

Physics  is  essentially  the  science  of  Energy,  and  aside  from  a  compara- 
tively brief  discussion  of  the  properties  of  Matter,  considers  the  latter 
only  as  it  is  associated  with  or  becomes  "the  vehicle  of  Energy." 

Chemistry  has  been  termed  the  science  of  matter.  This  is  due  to  the 
fact  that  heretofore  but  little  time  has  been  devoted  to  energy  considera- 
tions. Recent  years,  however,  have  witnessed  a  marked  change  in  the 
methods  employed  in  the  study  of  Chemistry.  More  emphasis  is  being 
placed  upon  the  role  of  energy  in  this  science.  With  this  change,  the 
artificial  line  of  demarcation,  separating  Physics  and  Chemistry,  has 
practically  disappeared. 

MATTER. 

If  we  consider  the  evidence  furnished  by  scientific  investigation,  it  is 
difficult  to  give  a  satisfactory  and  final  definition  of  matter.  It  is  better 
described  by  its  properties,  although  it  is  somewhat  evasively  defined  as 
anything  which  occupies  space  and  appeals  to  the  senses.  From  a 

*NOTE: — A  few  scientists  maintain  that  there  is  no  such  thing  as  "matter,"  that  it 
is  but  an  energy  manifestation;  others  hold  that  neither  matter  nor  energy  is  an  ob- 
jective reality,  but  merely  a  product  of  thought. 

17 


!8  EXPERIMENTAL  CHEMISTRY. 

chemical  point  of  view,  it  has  been  described  as  anything  which  possesses 
weight.  Science  seems  unable  as  yet  to  predicate  what  matter  is.  Ex- 
periment has  revealed  that  "the  total  mass  (quantity  of  matter)  of  any 
system  is  not  altered  by  any  process  which  may  take  place  within  that 
system." — Ostwald.  This  is  often  designated  as  the  "law  of  the 
conservation  of  matter."  (It  is  frequently  desirable  to  differentiate  the 
changes  occurring  within  a  body  or  a  "set  of  materials"  or  a  group  of 
bodies  from  those  changes  which  may  take  place  in  its  "surroundings" 
To  avoid  this  cumbrous  expression,  science  has  introduced  the  word 
"  system"  to  designate  this  assemblage  or  arrangement  of  bodies  considered 
as  being  insulated  from  its  environment.) 

Definite  portions  of  matter  are  called  "bodies"  or  "specimens." 
Different  kinds  of  matter  are  called  "substances."  Bodies  may  be 
homogeneous  or  heterogeneous  accordingly  as  they  are  made  up  of  visibly  * 
like  or  unlike  parts. 

Matter  appeals  to  the  senses  because  of  its  attributes  or  properties 
which  are  characteristic  of  a  body  or  a  substance. 

By  the  word  "property,"  the  idea  of  a  peculiar  quality  of  a  thing 
is  conveyed;-  - I>  .implies  that,  which  is  essential  and  inherent  in  a  thing, 
as  sweetness' Is  a'  property;  of  :sugar.  "The  properties  of  an  object  are 
all  the  relations  in,  whjciu  it  ,can  be(  made  to  appeal  to  the  senses." — 
Ostwald.//-  •  */•  ;„•[  { r ;  '  J  f  r ,\  f'-,  , r ; 

The  identity  of  a  body  or  a  substance  depends  upon  a  definite  assem- 
blage of  properties. 

The  matter  of  the  universe  is  continually  undergoing  a  change,  i.e., 
the  properties  are  being  altered  in  degree  or  completely  changed.  Any 
directly  observed  change  taking  place  in  matter  is  called  a  phenomenon. 

If  the  change  alters  the  properties  of  a  substance  but  temporarily,  it  is 
said  to  be  a  physical  phenomenon;  if  the  properties  are  changed  per- 
manently, i.e.,  the  substance  loses  its  original  identity,  the  observed 
change  is  called  a  chemical  phenomenon. 

Those  properties  which  are  exhibited  by  physical  phenomena  are 
called  physical  properties,  and  those  which  require  chemical  phenomena 
for  their  exhibition  are  called  chemical  properties. 

(Physics  has  sometimes  been  defined  as  the  study  of  physical  phe- 
nomena, and  Chemistry  as  the  study  of  chemical  phenomena.) 

Physical  properties  as  a  class  are  further  subdivided  into  two  groups, 
general  and  specific  properties.  General  properties  are  those  which 
are  possessed  in  common  by  all  kinds  and  conditions  of  matter.  Specific 
properties  are  those  which  are  characteristic  of  a  particular  kind  of 
matter  (a  substance),  yet  common  to  a  given  species  of  matter.  Note 
that  "bodies"  cannot  always  be  differentiated  from  one  another  by  the 
sole  use  of  specific  properties,  for  if  they  be  "bodies"  or  "specimens" 
of  one  kind  of  matter,  say  sulphur,  then  all  the  bodies  will  possess  in 
common  the  specific  properties.  Observe  that  such  attributes  as  size 

*  Colloidal  solutions  represent  an  exception  to  this  rule  if  the  word  "  visibly  "  is 
used  in  the  ordinary  sense  of  the  term. 


FUNDAMENTAL   CONCEPTS.  1 9 

and  form  assist  in  differentiating  "bodies,"  yet  said  attributes  do  not 
enter  into  our  concept  of  the  substance  of  which  the  body  is  composed. 

Experiment  I. — Matter,  Bodies  and  Substances. 

(Record  all  data  in  your  notebook.) 

Place  two  small  pieces  of  each  of  the  following  upon  the  top  of  your 
laboratory  desk — zinc,  iron,  sulphur  and  glass.  How  many  "bodies" 
are  represented?  How  many  substances?  How  did  you  differentiate 
the  bodies?  The  substances?  Name  the  properties  which  you  made 
use  of  in  each  case. 

A  piece  of  sulphur  (one  cubic  cm.),  under  the  ordinary  conditions  of 
the  laboratory  possessed  the  following  physical  properties:  color,  yellow; 
odor,  practically  none;  hardness,  2.5;  melting  point,  115°  C.;  boiling 
point,  448°  C.;  specific  gravity,  2;  specific  volume,  .5;  specific  heat,  1.8; 
insoluble  in  water;  soluble  in  carbon  bisulphide;  crystalline  structure; 
opaque;  poor  conductor;  weight,  2  grams;  form,  cubical;  heat  capacity, 
1.8  x  2;  inertia;  indestructibility;  extension;  porosity;  impenetrability. 

Do  all  of  the  properties  enumerated  above  belong  to  all  bodies? 
To  all  substances  ?  Classify  all  of  the  above  properties  under  the  three 
heads — "general,"  "specific,"  and  "body"  properties. 

Which  class,  or  classes  of  properties  enter  into  your  concept  of  matter? 
Bodies?  Substances? 

Can  you  differentiate  bodies  by  the  use  of  such  attributes  as  weight, 
form  and  size?  Try  to  do  so  by  placing  the  two  pieces  of  sulphur  side 
by  side.  Can  you  conceive  of  any  exception  to  your  conclusion  ?  If  so, 
explain. 

PHYSICAL    PROPERTIES. 
Experiment  II. — Determination  of  Specific  Gravity.     (Quant.) 

(a)  Clean  and  thoroughly  dry  a  50  cm. 3  Erlenmeyer  flask;  weigh  it 
accurately  and  record  weight  of  same  in  your  notebook;  from  a  burette 
containing  distilled  water  at  the  temperature  of  the  laboratory,  measure 
into  flask  10  cm. 3  of  the  water.     (Recall  former  instructions  relative  to 
meniscus).     Weigh    as   quickly  as    possible.     Calculate  the  weight  of 
i  cm.3  of  water  at  the  temperature  of  the  laboratory.  The  weight  of  i  cm.3 
of   a  substance  is  called  its  density.     The  weight  of  the  volume  of  a 
substance  compared  with  the  weight  of  an  equal  volume  of  water  at  the 
temperature  of  its  maximum  density  (what  is  this  temperature  ?)  is  called 
its   specific  gravity,  i.e.,  the  ratio  is  given  this  name.     The  reciprocal 
of  density  or  specific  gravity  of  a  body,  or  the  volume  occupied  by  one 
gram,  is  called  its  specific  volume. 

(b)  (Quant.)     To    determine  the  specific  gravity  of   glass — suspend 
a  piece  of  silk  thread  from  the  hook  on  the  balance  and  adjust  balance 
so  that  pointer  makes  vibrations  of  equal  length  on  either  side  of  the  zero 
point;  tie  the  thread  to  the  piece  of  glass  so  that  when  a  beaker  of  glass 
is  placed  under  it,  the  solid  will  be  completely  immersed  in  the  water; 


20  EXPERIMENTAL  CHEMISTRY. 

weigh  the  suspended  glass,  first  in  the  air  and  then  when  immersed. 
Record  the  two  weights  in  your  notebook.  What  is  the  difference  in 
weight  ?  What  is  the  weight  of  the  volume  of  water  equal  to  the  volume  of 
the  glass?  What  is  the  specific  gravity  of  glass?  Its  density?  The 
volume  of  the  piece  of  glass?  Its  specific  volume? 

(c)  (Quant.)     Determine  the  specific  gravity  of  alcohol  by  means  of 
a  specific  gravity  bottle  or  an  Ostwald-Sprengel  pyknometer.     (Instruc- 
tions from  assistant.) 

(d)  (Quant.)     To  ascertain  the  specific  gravity  of  a  solid  when  only 
small  pieces  are  available — weigh  the  specific  gravity  bottle  (Richard's 
form  preferred)  when  empty;  when  filled  with  water  and  properly  ar- 
ranged; weigh  the  solid;  place  the  solid. in  the  bottle  and  add  water  until 
bottle  is  filled,  observing  all  precautions;  weigh.     Weight  of  contents? 
Of  solid  in  bottle?     Of  water  now  in  bottle?     Does  the  solid  displace 
its  own  volume  of  water?     What  is  the  weight  of  this  volume  of  displaced 
water?     Using  small  pieces  of  glass,  determine  its  specific  gravity. 

Experiment  III. — Compressibility  of  Gases. — Boyle's  Law.     (Quant.) 

To  determine  the  relation  between  volume  and  pressure.  Plot  graph 
of  results.  (Instructions  from  assistant.) 

Experiment  IV. — Expansion  of  Gases  With  Increase  of  Temperature.— 
Charles'  Law.  (Quant.)  (Instruction.) 

Experiment  V. — Melting  and  Boiling  Points.    (Quant.)    (Instructions.) 
See  Traube's  "Physico-chemical  Methods." 

ENERGY. 

A  body  may  possess  other  qualities  or  "conditions,"  such  as  motion, 
electrical  charges  or  temperature.  These  conditions  may  be  regarded 
as  energy  relations  of  matter.  They  do  not  enter  into  our  concept  of  a 
substance,  yet  their  influence  upon  its  properties  is  very  marked  in  some 
cases. 

Energy  is  the  capacity  for  work.  "  It  is  the  essential  thing  in  the 
universe." — Richards,  T.  W.  By  some  it  has  been  called  the  funda- 
mental property  of  the  conceived  entity,  called  "matter."  This  is  prob- 
ably due  to  the  fact,  as  science  holds,  that  matter  is  always  associated 
with  more  or  less  energy.  The  idea  of  considering  energy  as  a  property, 
thereby  subordinating  it  to  matter,  is  severely  criticised  by  many,  who 
contend  that  it  should  be  placed  on  the  same  plane  with  matter.  The 
changes  in  properties  and  the  power  to  produce  them  are,  therefore,  con- 
ceived to  arise  not  from  a  number  of  distinct  entities,  but  from  a  single 
one,  which  is  capable,  however,  of  manifesting  itself  in  a  variety  of  differ- 
ent ways.  "That  which  gives  rise  to  the  changes  in  the  properties  of 
bodies  and  to  the  power  to  produce  such  changes,  is  called  energy." — 
Noyes,  A.  A. 

The  familiar  forms  of  energy  are,  heat  energy,  light  energy,  electrical 
energy,  gravitational  energy,  kinetic  energy,  etc.  Careful  experimentation 
covering  a  long  period  of  time  has  shown  that  equal  amounts  of  the  various 


FUNDAMENTAL    CONCEPTS.  21 

forms  of  energy  will  produce,  on  transformation,  equal  amounts  of  heat. 
This  has  led  to  the  enunciation  of  the  law  called  "conservation  oj  energy" 
which  may  be  stated  as  follows: 

In  a  limited  system,  regardless  oj  the  transformations  which  may  take 
place  within  the  system,  the  total  amount  oj  energy  is  not  altered. 

Julius  Robert  Mayer,  in  1842,  was  the  first  to  formulate  the  above 
law.  It  is  sometimes  called  the  First  Law  of  Energetics. 

The  importance  of  the  role  of  energy  in  Chemistry  gradually  becomes 
more  clear  as  the  individual  investigates.  It  is  generally  admitted  that 
"the  existence  of  matter  is  inferred  only  from  its  various  energy  mani- 
festations." This  leads  to  the  conclusion,  that  at  least  so  far  as  our  ex- 
perience is  concerned,  matter  and  energy  are  inseparably  connected; 
that  a  "  body"  as  we  observe  it,  represents  so  much  of  the  entity,  "  matter," 
associated  with  more  or  less  of  the  entity,  "  energy."  Matter  *  never  exists 
alone,  i.e.,  isolated  from  energy,  and  science  is  quite  sure  that  if  the  energy 
were  removed  from  a  piece  of  iron,  that  its  properties  would  be  very 
different.  A  body  or  an  object  possesses  in  addition  to  matter  a  certain 
"  energy  content,"  which  determines  its  properties.  This  form  of  energy, 
we  call  internal  energy.  (This  does  not  include  kinetic  energy.) 

Experiment  I. — Forms  of  Energy.     Transformation  of  Energy. 

(a)  Raise    the    object    called    a    "mortar"    from    your    desk.      Did 
this  require  energy?     Work?     What   kind  of   energy  did  you  exert? 
What  kind  of  energy  did  you  work  against  ?     What  became  of  the 
energy  you  expended?     Was  there  any  energy  destroyed  in  the  above 
operation  ?     Consider  yourself  and  the  mortar  a  body  with  its  capacity 
for  gravitational  energy,  as  constituting  a  "system;"  did  the  system  gain 
or  lose  any  energy?     Was  there  a  transformation  of  energy  during  the 
process?     State  specifically  the  transformation. 

(b)  Procure  a  hard  rubber  or  glass  rod  and  a  piece  of  flannel  from  the 
assistant;  observe  that  it  is  somewhat  "cold"  to  the  touch  and  that  if 
touched  to  pieces  of  thin  paper  it  does  not  disturb  them;  now  vigorously 
rub  the  rod  with  the  flannel;  hold  the  rod  close  to  the  pieces  of  paper; 
what  is  the  effect?     Touch  the  rod  and  note  whether  it  seems  to  have 
changed  in  temperature.     The  "rubbing"  required  the  expenditure  of 
what  kind  of  energy  ?     It  was  transformed  into  what  two  forms  of  energy  ? 
What  bodies  constitute  our  system  in  this  case  ?     Did  the  system  gain  or 
lose  energy?     Explain.     Are  chemical  or  physical  phenomena  involved 
in  this  experiment?     Give  reason  for  your  answer. 

(c)  Ask   the   assistant  for  a  d'Arsonval  galvanometer,  and  pieces  of 
iron  and  copper  wires  about  20  cm.  long.     Join  the  iron  and  copper  wires 
at  one  end  by  twisting  them  firmly;  attach  the  free  ends  of  the  wires  to  the 
two  binding  posts  of  the  galvanometer;  take  the  reading.     (Instructions.) 
Light  a  match  and  apply  the  flame  to  the  juncture  (twist)  of  the  two 
wires;  take  the  reading  while  applying  the  heat;  this  latter  form  of  energy 

*  There  is  much  data  at  hand  which  points  to  the  conclusion  that  "matter"  is 
nothing  but  energy. 


22 


EXPERIMENTAL  CHEMISTRY. 


is  partially  transformed  into  what  two  forms  of  energy  revealed  by  the 
conduct  of  the  galvanometer? 

(d)  (Instructions.)  Into  a  cold  dilute  solution  of  hydrochloric  acid  dip 
a  strip  of  zinc  and  a  strip  of  platinum;  by  means  of  wires  connect  strips 
with  the  binding  posts  of  an  ammeter.  Result  ?  Place  your  hand  upon 
the  vessel  containing  the  acid  and  metals;  any  change  in  temperature? 
Is  the  zinc  being  consumed,  i.e.,  is  it  dissolving  in  the  acid?  Has  that 
portion  which  has  been  consumed  lost  its  identity,  i.e.,  do  you  see  anything 
in  the  acid  which  seems  to  possess  the  same  assemblage  of  properties  that 
zinc  possesses  ?  Those  phenomena  in  which  there  is  involved  a  change 
of  identity  of  a  substance  are  called  chemical  phenomena.  As  the  result 
of  these  chemical  phenomena,  has  energy  become 
available  for  transformation  into  heat  energy  and 
electrical  energy?  Enumerate  as  many  different 
forms  of  energy  as  you  can. 

From  a  study  of  the  foregoing  experiments  the 
student  should  readily  deduce  that  energy  manifests 
itself  in  various  forms,  many  of  which  are  familiar. 
A  form  of  energy  with  which  the  chemist  is  almost 
continuously  concerned  is  the  one  illustrated  in  the 
last  experiment — the  change  of  identity  suffered  by 
zinc  when  it  dissolved  in  the  hydrochloric  acid.     The 
zinc  strip  is  supposed  to  contain  a  given  amount  of 
and   under    definite    conditions,    a 
further   asserted   that   this  internal 


FIG.  i. 


zinc   matter 

definite  quantity  of  energy.  It  is 
energy  which  is  "  stored  up  "  in  the  zinc  substance  is  liberated,  at  least 
in  part,  as  "free  energy"  when  the  zinc  presses  into  solution,  and  that  it 
is  this  so  called  jree  energy  which  is  transformed  into  heat  and  electricity. 
This  relationship  of  matter  and  energy  will  receive  attention  as  we  ad- 
vance in  our  work. 

Experiment  II. — Relation  of  Energy  to  the  Properties  of  Matter. 

(Postulating  that  our  senses  can  respond  to  only  energy  manifestations, 
proceed  with  the  following  experiments.) 

(a)  Procure  a  cylinder  of  sulphur  (or  hard  rubber  or  glass  rod)  and  a 
piece  of  flannel  from  the  assistant;  examine  the  piece  of  sulphur;  why 
does  it  appeal  to  the  senses?  Do  these  properties  bear  any  relation  to 
energy?  (Answer  in  the  light  of  the  above  postulate.)  If  they  do  not, 
should  properties  alter  with  change  of  "  energy  content  ?"  Test  your 
conclusion  by  experiment;  hold  the  sulphur  rod  near  small  pieces  of 
thin  paper;  has  it  the  property  of  attracting  the  paper?  Rub  the  rod 
vigorously  with  the  flannel;  now  hold  rod  near  the  paper.  Is  there 
any  alteration  or  change  in  the  properties  of  the  rod  ?  Did  you  expend 
energy  in  rubbing?  If  so,  what  became  of  it?  Explain  in  terms  of 
energy  relations. 

Rub  the  rod  again  with  flannel,  supposedly  "storing  up"  an  electrical 


FUNDAMENTAL    CONCEPTS.  23 

charge  on  the  sulphur;  is  its  energy  content  thereby  changed?  Has 
this  changed  its  identity?  Is  it  possible  that  many  of  its  properties  are 
altered  in  so  small  degree  that  we  are  unable  to  detect  the  change  ?  Does 
an  alteration  in  small  degree  of  the  properties  occasion  a  loss  of  identity  ? 

If  a  small  change  in  the  energy  condition  produces  a  small  change  in 
the  properties  of  a  body  or  a  substance  as  in  the  case  of  sulphur,  might  we 
be  led  to  assume  that  great  energy  changes  will  be  followed  by  correspond- 
ingly great  alterations  in  properties  of  bodies,  perhaps  to  the  extent  of 
changing  the  identity  of  a  substance,  i.e.,  nearly  a  complete  alteration  of 
properties  ? 

(b)  Examine  the  filament  in  the  incandescent  lamp  over  your  desk; 
note  its  properties  as  best  you  can  by  looking  through  the  glass;  observe 
especially  its  color,  its  diameter  and  its  light-giving  properties.  In  view 
of  your  previous  experience,  would  you  say  that  the  filament  possesses 
a  definite  amount  of  internal  energy  ?  Does  the  filament  possess  a  defi- 
nite assemblage  of  properties  under  the  present  conditions  of  temperature  ? 

Turn  the  switch  permitting  the  electricity  (energy)  to  flow  through  the 
filament.  This  energy  is  transformed  into  what  two  forms  of  energy 
easily  observed?  Does  the  filament  contain  more  or  less  energy  now 
than  when  current  is  not  traversing  it  ?  Are  the  properties  of  the  fila- 
ment altered  when  the  current  is  flowing  through  it?  Does  it  lose 
its  identity  either  temporarily  or  permanently?  Cut  off  the  current,  re- 
storing the  filament  to  the  original  " energy  content,"  i.e.,  to  the  same 
conditions  of  energy.  Does  it  regain  its  original  properties  ?  If  so,  what 
kind  of  a  phenomenon  have  you  been  observing? 

As  the  result  of  this  experiment  and  those  which  have  preceded  it, 
would  you  feel  justified  in  affirming  that  there  is  a  causal  relation  between 
energy  and  the  properties  of  matter?  Should  the  energy  conditions  be 
stated  before  an  attempt  is  made  to  describe  the  properties  of  a  substance  ? 
Why? 

It  seems  quite  certain  that  there  is  a  very  intimate  relation  between  the 
properties  of  a  substance  and  the  amount  of  energy  associated  with  it, 
that  the  properties  vary  in  degree  as  the  amount  of  internal  energy 
varies.  It  is  quite  likely,  if  we  could  remove  the  greater  portion  of  the 
internal  energy  possessed  by  sulphur  without  resorting  to  chemical  proc- 
esses, that  we  should  have  a  substance  whose  properties  would  be  very 
different  from  those  of  sulphur. 

The  above  experiment  with  the  filament  of  the  incandescent  lamp 
emphasizes  the  fact  that  when  we  speak  of  a  substance  possessing  a  par- 
ticular identity,  i.e.,  a  definite  set  of  properties,  we  mean  under  certain  de- 
fined conditions  of  energy. 

Experiment  III.— Relation  of  Energy  to  Phase  of  Substance. 

Examine  several  small  pieces  of  ice,  it  is  said  to  be  water  in  the  solid 
state  or  phase.  Has  it  a  definite  identity  ?  Place  a  few  small  pieces  of  the 
ice  in  a  test  tube  and  gently  warm  the  tube  and  its  contents  in  a  flame 


24  EXPERIMENTAL  CHEMISTRY. 

(heat  energy).  As  long  as  ice  and  water  are  together  in  the  test  tube  the 
temperature  of  the  mixture  does  not  alter.  What  becomes  of  the  energy 
that  is  being  contributed  by  the  flame  ?  When  all  of  the  ice  has  melted 
remove  the  test  tube  from  the  flame;  you  now  have  water  in  the  liquid 
phase.  Does  this  same  weight  of  water  contain  more  or  less  internal 
energy  than  the  equal  weight  of  ice  ?  Have  the  properties  been  altered  ? 
If  you  had  never  seen  water  in  the  liquid  phase,  would  you  have  recognized 
it  by  virtue  of  your  knowledge  of  the  properties  of  ice  ?  Has  the  identity  of 
the  ice  been  lost  ?  Introduce  the  tube  and  its  contents  into  the  flame  again, 
and  heat  until  all  of  the  latter  has  passed  into  the  gaseous  phase.  Can 
you  see  steam  ?  Does  the  steam  possess  more  or  less  internal  energy  than 
the  water?  Are  its  properties  different  from  those  of  water,  i.e.,  has  the 
steam  an  identity  of  its  own?  Are  these  phenomena  of  a  chemical  or 
physical  nature?  Would  you  say  that  the  energy  content  determines, 
at  least  in  a  large  measure,  the  phase  or  state  of  aggregation  of  water? 
Which  phase  of  a  substance  has  the  maximum  amount  of  internal  energy? 

A  solid  is  frequently  said  to  possess  "form  energy,"  by  which  expres- 
sion we  understand  that  sufficient  attraction  (cohesive  energy)  exists 
among  its  particles  to  give  it  rigidity  and  form.  Because  of  these  attri- 
butes due  to  its  energy  "condition,"  a  solid  is  sometimes  denned  as  pos- 
sessing the  properties  of  specific  volume  and  form.  A  liquid,  possessing 
less  effective  jorm  energy,  has  not  the  property  of  a  definite  form,  but 
that  of  specific  volume  only.  A  gas  has  neither  the  specific  property  of 
form  nor  definite  volume;  it  possesses  relatively  the  minimum  of  effective 
form  energy. 

When  the  ice  in  the  above  experiment  was  heated,  sufficient  energy  was 
imparted  to  it  to  overcome  the  effects  of  the  form  energy,  at  least  in  part, 
with  the  result  that  the  solid  was  changed  into  the  liquid  "  state  of  aggre- 
gation." "This  latter  expression  came  into  use  because  it  is  commonly 
assumed  that  in  the  different  states,  the  particles  of  which  a  substance  is 
composed,  are  differently  aggregated."  Heat  energy,  when  imparted  to 
these  particles,  acts  like  a  " repellent  force"  overcoming  the  form  energy 
and  causing  a  change  of  state  of  the  substance.  A  similar  course  of 
reasoning  is  followed  in  explaining  a  change  from  the  liquid  to  the  gaseous 
state. 

If  form  energy  is  dominant  in  a  substance,  then  the  solid  state  results; 
if  heat  energy  is  dominant,  then  the  gaseous  state  prevails. 

The  energy  content  of  a  substance  under  definite  pressure  determines 
its  physical  state  or  phase.  Considerable  confusion  has  grown  out  of  a 
careless  use  of  these  last  two  words. 

PHASES  AND  STATES  OF  MATTER. 

If  we  consider  ice,  water  and  water  vapor  (gas)  as  constituting  a  system, 
then  the  three  physically  distinct  parts  of  this  system  represent  three 
phases,  also  three  states — solid,  liquid  and  gaseous.  What  is  the  dis- 
tinction between  "phase"  and  "state?"  Ordinarily  we  consider  but 


FUNDAMENTAL    CONCEPTS.  25 

three  states  of  matter,  namely,  solid,  liquid  and  gaseous,  no  difference 
whether  we  are  speaking  of  and  comparing  various  substances  or  just  one. 
The  word  phase  is  usually  defined  as  a  homogeneous  aggregation  or  mass. 
Homogeneous  means  "like,"  i.e.,  possessing  same  properties;  heterogene- 
ous expresses  the  idea  of  "unlike."  If  kerosene  were  poured  upon  a 
little  water  in  a  test  tube  and  two  distinct  layers  were  seen,  how  many 
phases  would  be  present  ?  How  many  states  ?  It  is  readily  seen  that  we 
have  a  mass  of  -water  and  kerosene  which  is  in  the  liquid  "state,"  but 
that  we  have  two  phases,  i.e.,  two  masses  which  are  each  homogeneous. 
Aggregations  of  matter  which  are  visibly  different  or  can  be  mechanically 
separated  from  one  another  are  called  phases.  Examine  a  piece  of  granite 
and  there  will  be  visible  three  kinds  of  matter  physically  different,  and 
which  can  be  separated  by  a  mechanical  process,  hence  we  have  three 
phases  and  but  one  state  (solid),  represented.  The  granite  is  not  a  homo- 
geneous substance;  ice  is.  The  student  should  remember  that  the  num- 
ber of  different  substances  present  does  not  necessarily  determine  the 
number  of  phases.  When  you  make  a  solution  of  salt  and  water  you  have 
but  one  phase,  because  the  entire  mass  is  homogeneous.  Consider  a  vol- 
ume of  air;  it  is  composed  of  many  substances,  nitrogen  and  oxygen,  chiefly, 
yet  there  is  but  one  phase,  the  gaseous. 

Homogeneity  must  always  be  the  test.  The  terms  phase,  homogene- 
ous and  heterogeneous  are  customarily  restricted  to  usage  in  the  physical 
sense. 

Experiment  IV.— A  Study  of  Phases  of  Matter. 

Place  5  cm.3  of  water  in  a  test  tube;  how  many  "states?"  Phases? 
In  answering  this  question  ignore  air  and  watery  vapor  above  surface  of 
water.  Add  3  cm.3  of  alcohol;  shake  tube  well.  You  have  a  mass  of 
alcohol  and  water;  is  it  homogeneous?  How  many  "states"  present? 
Phases  ? 

Repeat  above,  using  kerosene  instead  of  alcohol.  Answrer  above  ques- 
tions in  order. 

Place  5  cm.3  of  water  in  a  test  tube;  add  2  grams  of  salt;  shake  well  and 
heat  gently  for  a  few  seconds.  How  many  phases?  States?  Add  5 
grams  of  salt  to  the  solution;  treat  as  before.  Do  you  find  that  all  of  the 
salt  will  not  go  into  solution?  How  many  states?  Ho\v  many  phases? 

From  the  above  it  is  evident  that  the  test  of  "  homogeneity"  (used  in 
the  physical  sense),  applied  to  a  mass — solid  or  liquid — will  usually  answer 
the  question  as  to  the  number  of  phases  present.  The  physically  dis- 
tinct parts  of  a  system  represent  the  phases  of  it.  There  is  but  one  phase 
recognized  in  the  gaseous  "state,"  but  many  in  the  solid  and  liquid  states. 


CHAPTER  II. 
CHARACTERISTICS  OF  PHYSICAL  AND  CHEMICAL  CHANGES. 

Those  phenomena  in  which  matter  undergoes  a  temporary  change  of 
identity,  regaining  its  original  identity  when  the  original  conditions  of 
energy  are  again  established,  are  called  physical  changes.  When  the 
change  in  identity  is  permanent,  it  is  known  as  a  chemical  change.  As  a 
rule,  chemical  changes  are  accompanied  by  physical  changes.  Chemical 
changes  are  defined  as  those  changes  which  affect  the  identity,  and 
produce  alterations  in  the  substance  under  defined  conditions  of  energy. 
Chemistry  is  primarily  the  science  of  these  chemical  changes  and  their 
attendant  phenomena. 

Experiment  I. — Physical  and  Chemical  Changes. 

(a)  Examine  the  properties  of  a  piece  of  platinum  wire  which  has 
been  sealed  into  the  end  of  a  glass  rod;  note  its  physical  properties  so 
carefully  that  you   become  sure  of  its  identity.     Has  it   a  particular 
group  of  properties  under  the  prevailing  conditions  of  energy?     Now 
hold  the  platinum  wire  in  the  Bunsen  flame  for  a  few  seconds.     Note 
its  properties   wrhile  in   the   flame.     Are   they  altered?     Remove  wire 
from  flame,  does  it  continue  to  glow  ?     Does  it  emit  light  now  ?     When 
you  are  confident  that  wire  is  under  the  original  energy  conditions,  ex- 
amine its  properties.     Has  it  lost  its  original  identity  ?     What  kind  of  a 
change  did  the  platinum  wire  undergo  ?     Define. 

(b)  Note  the  properties  of  a  piece  of  magnesium  " ribbon;"  by  means 
of  a  pair  of  pinchers,  hold  the  ribbon  in  the  flame  for  a  few  seconds  and 
then  remove  it;  does  it  give  out  light  and  heat  after  it  is  removed  from 
the  flame  ?     Examine  what  is  left  of  the  ribbon,  i.e.,  the  white  powder, 
when  it  is  cooled  to  the  original   external   energy  conditions;   has  its 
identity  been  changed?     What  kind  of  a  change  did  the  magnesium 
ribbon  undergo?     Define. 

Magnesium  +  Oxygen  — *  Magnesium  Oxide. 

In  the  above  experiment  it  must  be  evident  to  the  student  that  more 
energy  is  yielded  by  the  burning  of  the  magnesium  ribbon  than  was 
supplied  to  it  by  the  flame.  Is  it  possible  that  when  it  combined  with 
oxygen,  they  both  gave  up  a  quantity  of  energy  which  was  transformed 
into  heat  and  light  ?  If  so,  would  the  energy  content  of  magnesium 
oxide  be  less  than  the  sum  of  the  respective  energy  contents  of  magnesium 
and  oxygen? 

What  kind  of  changes  are  illustrated  by  the  rusting  of  iron,  a  falling 
ball,  burning  coal,  melting  of  ice,  and  the  souring  of  milk  ? 

26 


CHARACTERISTICS    OF    PHYSICAL    AND    CHEMICAL    CHANGES.  27 

Experiment  II.— (L.  T.)     Electrolysis  of  Water.— Change  of  Identity. 

Make  a  mixture  of  acid  and  water  in  the  ratio  of  i  to  20.  (Pour  the  acid 
into  the  water  slowly.)  Fill  a  Hoffman  apparatus  (Figs.  2  and  3)  for  the 
electrolysis  of  water  with  the  above  mixture.  To  each  of  the  stop  cocks 
attach  a  glass  delivery  tube  by  means  of  pieces  of  rubber  tubing,  so  that 
they  will  dip  into  a  vessel  of  water,  a  pneumatic  trough  preferably.  Keep 
stop  cocks  closed  until  at  least  15  cm.3  of  gas  has  collected  in  the  tube 
which  contains  the  lesser  volume  of  gas.  What  is  the  volume  in  the 
other  tube  (approximately)  ?  Now  collect  in  separate  test  tubes  these 


FIG.  2.— Voltameter. 


FIG.  3. 


respective  volumes  of  gases,  or  portions  of  them,  by  "water  displace- 
ment," i.e.,  fill  a  test  tube  with  water  and  invert  it  with  mouth  under 
water.  Open  stop  cock  slowly  and  gas  will  be  forced  through  delivery 
tube,  displacing  water  in  test  tube.  Place  your  thumb  over  the  test  tube 
mouth  under  water  and  bring  its  mouth  upward  to  a  burning  match. 
What  happened?  Repeat  this  operation  with  the  other  tube  of  gas. 
Using  a  glowing  splinter  repeat  both  operations. 

The  larger  volume  of  gas  was  hydrogen;  the  other  gas  was  oxygen 
In  which  gas  did  the  splinter  burn  most  vigorously  ? 

After  thus  examining  the  properties  of  these  two  gases,  would  you  say 
that  the  water  has  lost  its  identity? 


28  EXPERIMENTAL  CHEMISTRY. 

Water  — >  Hydrogen  +  Oxygen. 

Water  is  obviously  a  more  complex  substance  than  either  hydrogen 
or  oxygen.  The  water  has  undergone  a  chemical  change  known  as 
decomposition  or  analysis. 

Experiment  III. — Effect  of  Heat  on  Mercuric  Oxide. 

By  means  of  a  folded  piece  of  paper  introduce  a  little  mercuric  oxide 

(red  precipitate)  into  a  dry  test  tube;  determine  combined  weight  of 
tube  and  powder  by  means  of  chemical  balance; 
note  properties  of  the  red  powder;  heat  tube 
gently.  Does  color  of  powder  change  ?  Hold  a 
glowing  splinter  down  in  the  mouth  of  tube 
while  heating.  Results?  Do  you  recall  having 
worked  with  a  gas  which  revealed  similar  proper- 
ties ?  Its  name  ?  Examine  the  sides  of  the  tube; 
do  they  show  a  metallic  luster?  What  is  it ?  In 
view  of  above  phenomena  what  kind  of  a  change 
would  you  say  you  have  observed  ? 

FIG. 4. — (Smith andKeller).  As  the  powder  in  the  tube  cools  does  it  tend 
to  regain  its  original  color?  Is  this  action  the 

reverse  of  the  one  which  takes  place  when  tube  is  being  heated  ?     Weigh 

tube  and  contents.     Result  ? 

Mercuric  oxide  T*  Mercury  +  Oxygen. 

Of  the  three  substances  indicated  in  the  above  equation  which  is  the 
most  complex  ? 

The  mercuric  oxide  has  undergone  what  particular  kind  of  chemical 
change  ? 

Definite  conditions  of  energy  seem  to  be  necessary  to  secure  the  de- 
composition of  mercuric  oxide.  Is  it  possible  that  the  energy  con- 
tributed by  the  flame  is  stored  up  in  the  mercury  and  the  oxygen  ? 

Experiment  IV.— Effect  of  Heating  Copper  in  Air. 

Clean  and  dry  the  small  porcelain  crucible  which  you  will  find  in  the 
drawer;  introduce  into  it  about  i  gram  of  powdered  copper;  by  means 
of  chemical  balance,  find  total  weight  of  copper  and  crucible;  place 
crucible  upon  a  pipe-stem  triangle  (Fig.  4)  and  heat  with  Bunsen  flame 
for  15  to  20  minutes.  Is  there  any  change  in  color?  While  the  crucible 
and  contents  are  cooling  examine  the  material  upon  the  side  shelves, 
labeled  ''copper  oxide."  Does  it  bear  any  likeness  to  substance  in 
crucible?  Place  crucible  and  contents  upon  balance.  Have  they 
gained  or  lost  weight?  Explain.  Which  is  the  more  complex,  copper, 
or  copper  oxide  ? 

Copper  +  Oxygen  — *  Copper  Oxide. 


CHARACTERISTICS    OF    PHYSICAL    AND    CHEMICAL    CHANGES.  29 

Chemical  changes  like  the  one  above  involving  the  addition  of  sub- 
stances to  one  another  are  called  combinations  or  syntheses.  In  the 
experiment,  you  arranged  conditions  favorable  for  the  synthesis  of 
copper  oxide. 

Experiment  V. — Solution  Facilitates  Chemical  Action. 

(a)  Place  about  .5  grm.  of  tartaric  acid  in  a  mortar  and  pulverize  it; 
add  to  contents  of  mortar  .5  grm.  of  sodium  hydrogen  carbonate  (bicar- 
bonate of  soda);  note  that  no  chemical  action  takes  place;  pulverize  the 
mixture;  any  chemical  action?     Transfer  a  little  of  the  powder  to  a  test 
tube  and  add  water;  does  a  chemical  change  take  place? 

(b)  Add  5  cm.3  of  concentrated  hydrochloric  acid  to  a  test  tube;  add 
the  same  quantity  of  ammonium  hydroxide  to  another  tube;  hold  the 
mouths  of  the  test  tubes  near  to  each  other.     What  evidence  of  a  chemical 
change?     Explain.     (Both  of  the  above  reagents  are  gases  dissolved 
in  water.) 

Chemical  action  occurring,  in  which  one  or  more  substances  are  solids, 
must  be  largely  a  surface  phenomenon,  i.e.,  occurring  on  the  surface.  The 
extent  of  surface  will  determine  largely  the  rate  at  which  the  action 
proceeds.  Substances,  finely  divided  expose  more  surface  per  unit 
weight,  hence  give  more  intimate  contact  with  one  another.  Solutions 
of  substances  provide  a  convenient  method  for  securing  intimate  contact 
of  substances.  Gases  provide  an  ideal  condition  but  are  bulky  and 
inconvenient  to  manipulate. 

An  interesting  application  of  this  principle  of  determining  the  rate 
at  which  the  chemical  action  shall  proceed  by  determining  the  area 
of  surface  exposed,  is  utilized  in  the  manufacture  of  explosives.  The 
explosives  are  cast  into  cylindrical  sticks  and  the  surface  is  increased  by 
providing  a  number  of  longitudinal  holes  of  varying  diameters. 

Experiment  VI. — Relation  of  Energy  to  Chemical  Change. 

From  the  reagent  bottle  containing  silver  nitrate,  procure  i  or  2 
cm.3  of  the  solution  in  a  clean  test  tube.  Dip  the  end  of  the  glass 
stirring  rod  into  the  liquid  and,  using  it  like  ink,  write  your  name  upon  a 
piece  of  white  paper;  expose  the  writing  upon  paper  to  direct  sunlight 
for  half  an  hour — if  need  be  leave  it  until  next  laboratory  period.  Re- 
sults? Considering  your  method  of  procedure,  what  factor  produced 
the  chemical  change? 

EXERCISES. 

(Record  all  answers  in  notebook.) 

What  form  of  energy  caused  the  change  in  the  platinum  wire  ?  De- 
composition of  water?  Analysis  of  mercuric  oxide?  Assisted  in  the 
synthesis  of  copper  oxide?  Change  of  silver  nitrate?  Does  energy 
seem  to  be  as  closely  identified  with  chemical  changes  as  it  was  with 
physical  changes  ?  Has  each  one  of  the  above  changes  involved  time  or 


30  EXPERIMENTAL  CHEMISTRY. 

have  they  taken  place  "  instantaneously  ?  "  (The  word  "  instantaneously  " 
is  seldom  found  in  the  vocabulary  of  a  scientist — what  do  we  mean  by 
this  affirmation?) 

What  is  the  chief  distinction  from  the  standpoint  of  energy  consider- 
ations between  the  conduct  of  the  piece  of  the  magnesium  ribbon  under- 
going chemical  change  and  the  other  chemical  changes  enumerated 
above  ?  Was  energy  necessary  to  induce  the  change  ? 

Name  the  different  kinds  of  chemical  changes  which  you  have  studied 
to  date. 


CHAPTER  III. 
ELEMENTARY  SUBSTANCES,  MIXTURES  AND  COMPOUNDS. 

It  appears  from  the  foregoing  experiments  that  under  suitable  con- 
ditions of  energy,  matter  can  be  made  to  combine  with  other  kinds  of 
matter  with  little  difficulty,  producing  more  complex  substances,  but  it 
has  not  been  possible  to  continue  simplifying  matter  indefinitely.  It 
has  been  impossible  to  go  beyond  resolving  all  kinds  of  matter  into 
anything  more  simple  than  about  80  so-called  elementary  substances. 
These  simple  substances  have  resisted  every  effort  to  date  to  decom- 
pose them  into  anything  more  simple  or  elementary.  Examples  of 
these  elementary  substances  are,  platinum,  magnesium,  mercury,  oxygen, 
copper,  gold,  silver,  iron,  hydrogen.  Water  is  not  an  elementary  sub- 
stance as  you  were  able  to  decompose  it  into  two  simple  substances, 
namely,  hydrogen  and  oxygen.  Substances  which  can  be  resolved  into 
something  more  elementary  are  not  elementary  substances.  It  has  been 
suggested  (Richards,  T.  W.)  that  this  limit  to  convertibility  might  be 
called  the  "Second  Law  of  Matter." 

Experiment  I. — Examination  of  the  Elementary  Substances. 

(L.  T.)  Ask  the  assistant  to  place  specimens  of  the  80  elementary 
substances  on  the  table,  then  carefully  note  the  physical  properties  of 
the  various  substances.  How  many  are  in  the  liquid  state  ?  Gaseous 
state?  Solid  state?  Notice  flourine,  chlorine,  bromine,  iodine;  is  there 
apparently  any  gradation  in  properties  (physical)  ? 

Observe  those  which  have  a  metallic  lustre.  The  majority  of  these 
are  called  metals.  Record  the  names  of  those  which  do  not  have  this 
property — they  are  called  non-metals. 

Experiment  II. — Energy  Content  Affects  Properties  of  Elementary  Sub- 
stances. 

(L.  T.)  Compare  the  properties  of  yellow  phosphorous  with  those 
of  red  phosphorous.  Perform  such  experiments  as  the  assistant  suggests. 
Note:  Don't  handle  the  phosphorus  with  the  hands  or  take  it  into  the 
vicinity  oj  much  heat;  cut  it  under  water. 

The  above  is  a  chemical  change  due  to  change  in  energy  content, 
causing  a  new  internal  rearrangement.  Consult  lecture  notes. 

Experiment  III. — Mixtures — Heterogeneous  and  Homogeneous. 

(a)  Try  to  dissolve  a  small  quantity  of  flowers  of  sulphur  in  water. 
After  shaking  vigorously  how  many  phases  have  you  ?  Can  you  separate 


32  EXRERIMENTAL  CHEMISTRY. 

the  t\vo  substances  by  a  mechanical  process  like  filtration  (instructions) 
or  evaporation  ?     Did  each  substance  retain  its  original  properties  ? 

(b)  Dissolve  a  small  quantity  of  flowers  of  sulphur  in  carbon  disulphide. 
(This  latter  substance  is  very  inflammable  and  quite  poisonous;  keep  it 
away  from  free  flames.)  How  many  phases  ?  Pour  a  little  of  the  solution 
into  a  beaker;  set  it  in  the  open  window  for  15  to  20  minutes  or  longer. 
Results?  Does  this  prove  that  you  had  a  mixture?  Why?  What 
kind  of  a  mixture? 

A  Mixture  is  an  aggregate  of  substances  in  which  each  substance 
retains  its  characteristic  properties.  The  various  substances  of  a  mix- 
ture are  called  "components."  One  of  the  best  tests  of  a  mixture  is,  can 
the  substances  be  separated  by  mechanical  processes. 

Experiment  IV. — Mixtures  and  Compounds. 

Mix  thoroughly  in  a  mortar  5  grams  of  iron  filings  and  an  equal  weight 
of  sulphur.  Place  a  portion  of  the  contents  of  mortar  on  a  sheet  of  paper 
and  pass  a  magnet  near  it.  Note  that  the  iron  is  withdrawn  leaving  the 
sulphur. 

Introduce  a  small  portion  of  the  contents  of  mortar  into  a  test  tube  and 
add  5  or  6  cm.3  of  carbon  disulphide;  shake  tube  vigorously;  does  the 
sulphur  dissolve  and  leave  the  iron  in  the  tube  ?  Pour  a  portion  of  the 
liquid  upon  a  watch  glass;  examine  after  15  or  20  minutes.  Results? 
What  particular  name  would  you  apply  to  the  contents  of  the  mortar  ? 

Now  take  the  portion  that  remains  in  the  mortar,  and  by  means  of  a 
folded  paper  introduce  it  into  a  test  tube  until  tube  is  half  filled;  heat 
test  tube  to  redness.  When  the  iron  and  sulphur  become  sufficiently 
hot  they  combine  and  the  mass  glows  brightly  although  tube  is  taken 
out  of  the  flame.  (Source  of  energy?)  Cool  and  break  the  tube. 
Examine  portions  of  contents  with  magnet.  Are  its  properties  the  same 
as  those  of  either  iron  or  sulphur?  What  kind  of  a  change  has  taken 
place  ?  The  new  substance  formed  is  called  a  compound.  Did  the 
action  continue  spontaneously  when  the  flame  was  removed  ?  Was 
heat  and  light  liberated  as  the  result  of  the  chemical  action  ? 

Iron  +  Sulphur  — »  Iron  Sulphide. 

When  elementary  substances  like  iron  and  sulphur  combine  to  form  a 
new  substance  \vhose  properties  are  characteristic,  and  in  wrhich  the 
properties  of  the  combining  substances  are  merged  and  lost,  we  say  a 
compound  has  been  formed.  The  various  substances  uniting  are  called 
the  ''constituents." 

All  matter  may,  for  purposes  of  convenience,  be  divided  into  three 
classes,  viz.:  elementary  substances,  mixtures,  and  compounds. 

Experiment  V. — Interaction  of  Compounds. 

Recall  the  interaction  of  hydrochloric  acid  and  ammonium  hydroxide. 


ELEMENTARY    SUBSTANCES,    MIXTURES    AND    COMPOUNDS.  33 

Experiment  VI. — Metathesis — A  Species  of  Chemical  Change. 

Place  a  few  pieces  of  potassium  iodide  in  a  mortar;  observe  its  color. 
Now  add  a  small  quantity  of  lead  nitrate;  observe  its  properties.  Is 
there  any  evidence  of  chemical  activity  ?  Pulverize  the  mass.  Is  there 
any  indication  that  a  chemical  change  has  taken  place  ?  What  is  it  ? 

This  is  a  case  of  "double  decomposition"  or  " metathesis."  As  shown 
by  the  following  equation,  it  is  a  reaction  in  which  the  "factors"  are 
represented  by  one  or  more  compounds  and  there  is  an  exchange  of 
position  between  two  or  more  substances. 

Potassium    Iodide  -f  Lead  Nitrate  — »  Lead  Iodide  -f  Potassium 

Nitrate. 
Copper  -f  Mercuric  Oxide  — -  Mercury  -f  Copper  Oxide. 

Metathesis  is  defined  as  a  reaction  in  which  a  substance  is  transferred 
from  one  compound  to  another. 

Substances  which  have  a  chemical  effect  on  one  another  are  said  to 
react;  the  chemical  change  is  called  a  reaction;  the  individual  substances 
which  participate  in  the  reaction  are  called  reagents  or  "factors;" 
the  new  substances  formed  are  called  the  "products."  The  properties 
which  enable  substances  to  react  on  one  another  with  varying  degrees 
of  activity  are  among  the  more  important  of  what  are  known  as  chemical 
properties. 


CHAPTER  IV. 
NOTE  ON  THE  ENERGETICS  OF  CHEMISTRY. 

THEORY. 

The  importance  of  the  role  of  energy  in  its  relation  to  matter,  and 
physical  and  chemical  changes  has  been  obvious  in  all  of  the  preceding 
work.  The  student  has  become  familiar  with  some  of  the  forms  of 
energy  which  may  be  associated  with  matter;  viz.,  light  energy,  heat 
energy,  electrical  energy,  magnetic  energy  and  gravitational  energy. 
(In  all  the  discussions,  energy  of  motion  of  the  body  as  a  whole,  i.e. 
kinetic  energy,  is  excluded.)  The  truthfulness  of  the  First  Law  of  energy 
has  become  more  apparent.  It  has  been  possible  to  transform  energy 
from  one  form  to  another.  We  know  from  observation  that  heat  energy 
supplied  by  the  combustion  of  fuel,  has  driven  steam  engines,  which  in 
turn  have  driven  dynamoes  supplying  electrical  energy  which  has  finally 
been  converted  into  heat  energy  and  light  energy.  We  have  understood 
fairly  well  this  series  of  energy  transformations,  but  there  has  been  one 
part  which  has  not  been  quite  so  plain,  that  is,  just  how  and  why  was 
heat  energy  evolved  by  the  combustion  of  coal  in  the  air?  Why  was 
heat  energy  and  light  energy  developed  by  the  burning  of  magnesium 
ribbon  in  the  air,  and  why  should  the  combining  chemically  of  iron  and 
sulphur  manifest  similar  phenomena  ?  It  was  not  quite  plain  why  heat 
energy  and  electrical  energy  should  make  their  appearance  when  strips 
of  zinc  and  copper  were  properly  arranged  in  a  dilute  solution  of  sul- 
phuric acid,  although  it  may  have  been  explained  by  saying  that  it  was 
due  to  the  chemical  reaction  of  zinc  upon  the  acid.  In  all  of  these  cases 
we  failed  to  observe  the  presence  of  any  of  these  familiar  forms  of  energy 
in  or  upon  the  various  substances  before  the  chemical  action  occurred, 
yet  we  know  that  energy  was  evolved,  and  that  it  must  have  come  from 
some  form,  possibly  some  other  form  of  energy  in  the  body  of  system. 
(First  Law  of  Energy.) 

The  question  fairly  before  us  is,  can  we  harmonize  and  interpret  these 
phenomena  in  the  light  of  our  original  concept  as  to  the  relation  of  matter 
and  energy?  Can  we  reach  some  conclusion  in  regard  to  the  relation 
which  properties,  physical  and  chemical,  bear  to  the  " matter  content" 
or  " energy  content"  or  both,  of  various  substances? 

As  was  stated  in  the  beginning  of  the  work,  matter,  as  we  know  it,  is 
conceived  of  as  consisting  of  an  entity,  "matter,"  always  associated 
with  a  definite  quantity  of  another  entity,  known  as  "energy,"  under 
defined  conditions.  For  illustration — a  piece  of  the  substance  iron 
as  we  know  it,  is  held  to  contain  so  much  "iron  matter"  and  so  much 

34 


NOTE    ON    THE    ENERGETICS    OF    CHEMISTRY.  35 

"internal  energy,"  i.e.,  any  body  under  defined  conditions  has  a  definite 
" matter  content"  and  a  definite  "energy  content."  We  can  not  wholly 
separate  the  "energy"  from  the  "matter."  What  we  really  see  when 
observing  a  piece  of  any  substance,  is  a  particular  species  of  matter, 
simple  or  compound,  associated  with  a  definite  quantity  of  internal 
energy. 

That  a  very  intimate  relation  does  exist  between  the  above  two  enti- 
ties is  evidenced  by  the  experiments  which  have  been  performed.  An 
alteration  in  the  energy  content  of  a  body  or  a  system  produced  certain 
changes  in  the  physical  properties  of  the  substances  involved.  Many 
times  the  physical  properties  seemed  to  be  sort  of  functions  -of 
the  energy  content.  It  was  observed  also  that  the  energy  conditions 
play  a  very  prominent  part  in  inducing  chemical  changes.  It  has  been 
obvious  in  many  of  the  chemical  phenomena,  if  the  Law  of  Conversation 
of  Energy  is  universal  in  its  application,  that  in  addition  to  the  change 
of  matter,  there  has  been  not  only  a  material  altering  in  the  quantity 
of  the  internal  energy  of  the  bodies  or  systems  undergoing  such  change, 
but  the  energy  has  been  frequently  transformed. 

A  more  definite  conception  as  to  what  we  mean  by  the  word  "  internal 
energy"  or  "energy  content"  may  aid  in  an  interpretation  and  explana- 
tion of  the  mechanism  of  the  phenomena  all  of  which  at  present  are  not 
understood. 

By  the  expression  internal  energy,  is  meant  the  total  energy,  regardless 
of  form,  within  a  body  or  system.  (This  does  not  include  kinetic  energy.) 
In  addition  to  the  forms  of  energy  already  enumerated,  and  which  may 
be  associated  with  matter,  there  are  other  forms  even  more  uniformly 
associated  with  matter,  namely,  cohesion  energy,  disgregation  energy 
and  chemical  energy.  The  first  two  are  those  forms  of  energy  which  a 
body  "possesses  in  virtue  of  the  tendency  of  its  particles  to  approach 
and  to  recede  from  one  another  respectively."  In  solids  the  cohesion 
energy  is  relatively  greater  than  the  disgregation  energy;  the  reverse  is 
true  of  gases,  where  the  particles  tend  to  repel  one  another.  Elasticity 
of  matter,  surface  tension,  contraction  of  substances  on  cooling,  etc., 
all  are  manifestations  of  cohesion  energy.  When  a  body  undergoes 
change  of  state  by  the  application  of  heat  energy,  that  which  is  directly 
accomplished  is  the  supplying  of  sufficient  disgregation  energy  to  over- 
come, at  least  in  part,  the  cohesion  energy.  It  is  evident  that  the 
internal  energy  of  the  body  has  been  increased,  i.e.,  the  energy  content 
is  now  greater  than  it  was  before  the  change  of  state.  It  follows,  owing 
to  this  relation  between  cohesion  energy  and  disgregation  energy,  that 
if  this  body  after  change  of  state  were  placed  in  such  a  relation  that  it 
did  not  continue  to  receive  heat  energy,  it  would  give  out  heat  to  its  en- 
vironment and  the  internal  energy  would  be  diminished  in  amount.  Any 
change  in  the  relation  of  these  two  forms  of  energy  in  a  body  will  be  followed 
by  a  change  in  the  internal  energy,  i.e.,  heat  energy  will  either  be  absorbed 
or  evolved.  Physical  changes,  including  change  of  state,  receive  a 
comparatively  satisfactory  explanation  by  a  line  of  reasoning  simila 


36  EXPERIMENTAL  CHEMISTRY. 

to  the  foregoing,  but  not  so  with  the  majority  of  chemical  changes.  The 
question  might  be  asked  here,  what  "caused"  the  magnesium  and 
oxygen  of  the  air  to  react  on  one -another?  Were  cohesion  energy  and 
disgregation  energy  the  real  "cause"  of  the  reaction?  Science  says  no. 
Again,  was  the  energy  evolved  by  the  combination  of  magnesium  and 
oxygen  exactly  equal  to  the  diminution  of  either  cohesion  energy  or 
disgregation  energy  or  both,  in  the  system  composed  of  the  above  two 
substances  ?  It  is  true  that  these  two  forms  of  energy  may  have  been 
altered  in  degree,  but  there  seems  to  be  another  form  of  energy  asso- 
ciated with  matter  which  is  probably  the  main  source  of  energy  in  the 
above  example,  and  is  possibly  the  "cause"  of  the  reaction  of  substances. 
It  is  called  chemical  energy,  and  is  possessed  by  systems  in  virtue  of 
"this  tendency  of  the  substances  which  comprise  the  system,  to  undergo 
transformations  into  other  substances."  The  exact  nature  of  this  so- 
called  chemical  energy  has  not  been  determined,  but  it  may  be  compared 
with  potential  energy,  for  it  appears  to  depend  largely  if  not  altogether 
upon  the  relative  positions  of  matter  in  which  it  resides.  An  illustration 
of  this  is  a  mixture  of  hydrogen  and  oxygen  which  seems  to  possess 
potential  energy,  because  of  the  proximity  of  its  particles.  If  a  spark  is 
passed  through  the  mixture  a  chemical  change  results  with  the  trans- 
formation of  chemical  energy  into  heat  energy.  An  impulse  was  required 
to  induce  the  reaction.  (It  might  be  said  in  a  popular  way,  to  overcome 
"chemical  inertia.")  This  necessity  of  contributing  a  small  quantity 
of  energy,  relatively  negligible,  to  initiate  the  transformation  of  energy 
seems  to  be  characteristic  of  many  phenomena,  both  chemical  and 
physical.  We  observed  this  in  the  combustion  of  the  magnesium  ribbon 
and  in  securing  the  combination  of  iron  and  sulphur.  Afterwards  the 
reactions  proceeded  spontaneously.  A  ball  lying  at  the  top  of  an  in- 
clined plane,  requires  an  impulse  to  start  it. 

Chemical  energy  is  often  regarded  as  being  composed  of  two  factors, 
one  unknown,  and  the  other,  possibly,  chemical  affinity.  It  is  obvious 
that  chemical  energy  constitutes  but  a  part  of  the  internal  energy;  also 
let  it  be  noted  that  in  a  system  (other  conditions  being  the  same)  where 
there  is  considerable  tendency  of  substances  to  react  (affinity)  as  in  the 
case  of  hydrogen  and  oxygen,  or  magnesium  and  oxygen,  there  will  be 
represented  more  chemical  energy,  hence  more  internal  energy,  than 
in  the  case  of  a  system  where  there  is  little  or  no  tendency  of  the  substances 
to  react.  It  is  easily  seen,  then,  that  if  the  systems  are  alike  in  every 
other  respect,  that  in  the  former  case,  the  substances  will  combine  with  a 
greater  liberation  of  energy,  due  in  the  main  to  a  transformation  of  the 
chemical  energy.  It  can  not  be  deduced  from  this,  however,  that  the 
diminution  in  the  internal  energy  of  the  system,  which  can  be  measured, 
is  a  measure  of  the  chemical  energy  or  affinity  of  substances,  for  it  is 
possible  and  quite  likely  that  during  the  chemical  change  there  have 
been  changes  in  other  forms  of  energy,  such  as  the  cohesion  and  dis- 
gregation energies.  This  latter  change  would  also  produce  an  alteration 
in  the  heat-capacity  of  the  system  which  would  give  it  the  property  to 


NOTE    ON    THE    ENERGETICS    OF    CHEMISTRY.  37 

liberate  or  absorb  heat,  in  degree  beyond  that  which  the  system  originally 
possessed,  exclusive  of  its  chemical  energy.  Other  factors  enter  also  to 
prevent  ordinary  usage  of  the  "diminution  of  the  internal  energy"  of  a 
system  as  a  measure  of  the  " affinity"  of  substances  in  the  system,  but  a 
discussion  of  same  would  be  out  of  place  in  this  "note." 

The  amount  of  energy  available  for  work,  liberated  by  a  body  or  a 
system,  is  called  the  "free  energy"  (U),  while  that  which  is  not  available 
to  do  work  is  called  "bound  energy."  The  total  energy  (internal  energy) 
of  a  system  is  equal  to  the  sum  of  the  "free"  and  "bound"  energies. 
It  should  not  be  inferred  that  because  the  internal  energy  of  a  system 
is  large,  that  it  necessarily  follows  that  the  "free"  energy  content  is 
large.  The  chemical  energy  which  may  be  liberated  and  transformed 
into  free  energy  from  any  one  substance  will  differ  in  amount  as  it  reacts 
with  different  substances.  When  sixteen  grams  of  oxygen  combine 
with  hydrogen  to  form  steam  about  57,000  calories  of  heat  are  liberated; 
when  the  same  amount  of  oxygen  combines  with  solid  carbon  to  produce 
gaseous  carbon-dioxide,  about  48,500  calories  are  evolved. 

A  law  of  much  importance  in  chemistry  is, 

That  a  reaction  proceeds  in  the  direction  in  which  there  will  be  the 
greatest  diminution  oj  free  energy. 

This  "free  energy"  is  perhaps  most  frequently  liberated  as  "  heat 
energy." 

Another  generalization  with  which  every  student  of  chemistry  should 
be  acquainted,  and  which  will  enable  him  to  frequently  determine  the 
direction  of  a  given  reaction  when  placed  under  new  conditions  of  energy, 
is  known  as  Le  Chatelier's  Theorem.* 

"Any  change  in  the  /actors  of  equilibrium  from  outside,  is  followed  by 
a  reverse  change  within  the  system." 

APPLICATION    OF    THEORY. 

The  questions  proposed  regarding  the  explanation  of  various  chem- 
ical phenomena  remain  as  yet  unanswered.  What  interpretation  can  we 
offer  from  the  view  point  of  the  foregoing  theory  ?  Let  us  take  the  case 
of  the  combination  of  iron  and  sulphur.  These  are  elementary  sub- 
stances; they  are  not  elements.  The  word  "element"  is  restricted  to 
mean  "matter"  alone,  as  "iron  matter,"  or  "sulphur  matter."  It  is 
an  abstraction.  We  have  had  no  experience  with  the  "elements;"  that 
with  which  we  have  experience  are  the  two  "elements"  associated  with 
their  respective  and  appropriate  loads  of  energy.  If  the  energy  could 
be  removed,  what  would  they  look  like?  They  each  possess  a  given 
amount  of  internal  energy  (cohesion  energy,  disgregation  energy,  chem- 
ical energy).  The  two  simple  substances  were  ground  together  and  the 
mixture  placed  in  a  test  tube.  We  conceive  of  this  mixture  as  consti- 
tuting a  system.  Its  internal  energy  is  equal  to  the  sum  of  the  energy 
contents  of  iron  and  sulphur.  There  was  no  evidence  of  chemical 
action,  nor  anything  to  indicate  the  presence  of  chemical  energy,  yet 
when  a  small  quantity  of  heat  energy — an  impulse — was  imparted  to  it, 
*  Bancroft,  The  Phase  Rule. 


38  EXPERIMENTAL  CHEMISTRY. 

the  iron  and  sulphur  combined  with  the  evolution  of  considerable  light 
and  heat,  to  form  a  compound,  known  as  iron  sulphide.  The  heat  and 
light  were  the  result  of  the  transformation  of  chemical  energy  chiefly, 
with  a  possbility  that  a  portion  was  yielded  by  the  transformation  of 
the  cohesion  and  disgregation  energies  of  the  respective  substances  when 
they  passed  into  the  form  of  the  new  substances. 

The  compound  (iron  sulphide)  contains  the  two  "elements"  but  their 
energy  contents  are  very  different  from  what  they  were  originally,  and 
therefore  we  expect  the  properties  of  the  compounds  to  be  quite  differ- 
ent. Is  it  likely  that  if  we  were  able  to  return  to  the  "elements"  their 
original  and  respective  loads  of  energy  that  we  would  secure  a  return  of 
the  elementary  substances,  iron  and  sulphur?  Consider  the  decom- 
position of  mercuric  oxide.  It  will  be  recalled  that  the  reaction  con- 
tinued spontaneously.  In  most  cases  where  the  chemical  change  pro- 
ceeds as  above  to  a  completion,  the  free  internal  energy  of  the  system 
diminishes  by  being  transformed  into  other  forms  of  energy.  We  know 
that  in  the  combustion  of  carbon  (coal)  in  air  that  a  very  large  per 
cent  of  the  heat  of  combustion  is  convertible  into  free  energy  (external 
work). 

In  concluding  this  discussion,  it  might  be  well  to  give  an  example 
of  a  chemical  change  in  which  there  is  no  apparent  alteration  in  the 
"  matter,"  the  only  change  being  in  the  energy  content.  If  yellow  phos- 
phorous is  heated  in  a  closed  tube  out  of  contact  with  the  air  at  a  tem- 
perature of  250°  C.,  it  will  be  converted  into  the  "allotropic"  form 
known  as  red  phosphorous  with  an  evolution  of  heat. 

Yellow  Phosphorous.  Red  Phosphorous. 

Yellow  color,  Red  color, 

Poisonous,  Non-poisonous, 

Burns  at  low  temperature,  Burns  only  at  high  temperatures, 

Phosphorescent,  Non-phosphorescent, 

Soluble  in  C  S2>  Insoluble  in  C  S2, 

Garlic  odor,  No  odor. 

These  two  substances  so  different  in  properties  possess  in  common 
the  "element"  phosphorous,  but  the  internal  energy  of  each  is  different. 

Elementary  substances,  which  by  virtue  of  different  energy-contents, 
possess  different  properties,  are  called  allotropic.  Other  examples  are, 
graphite,  diamond  and  amorphous-carbon. 

THERMOCHEMISTRY. 

That  branch  of  chemistry  which  deals  with  the  thermal  (heat)  changes 
accompanying  chemical  reactions  is  known  as  Thermochemistry.  The 
chief  aim  of  Thermochemistry  is  to  determine  the  relative  "affinities"  of 
different  substances  by  the  principle  of  "the  development  of  the  great- 
est amount  of  heat."  This  as  ordinarily  used  furnishes  only  approxi- 
mate data. 

The  principles  of  thermochemistry  are  in  part  summarized  by  Bloxam 


NOTE    ON    THE    ENERGETICS    OF    CHEMISTRY.  39 

as  follows: — (i)  Every  chemical  change  is  accompanied  by  a  thermal 
change  which  is  a  constant  quantity.  (2)  The  thermal  change  occurring 
during  the  combination  of  elements  to  form  a  compound  is  called  the 
heat  of  formation*  If  heat  is  evolved  as  it  usually  is,  it  is  said  to  be  an 
exothermic  compound;  if  heat  is  absorbed,  endothermic.  Endothermic 
compounds  (usually  explosives)  contain  more  internal  energy  than  the 
constituents  originally  possessed,  and  are  unstable.  (3)  The  thermal 
change  occurring  during  the  decomposition  of  a  compound  is  called  the 
heat  of  decomposition.  (4)  The  heat  of  decomposition  is  identical  with, 
but  opposite  in  sign  to  the  heat  of  formation  of  a  compound. 

The  measurement  of  the  thermal  changes  is  accomplished  by  means 
of  a  calorimeter.  A  very  accurate  form  of  this  apparatus,  known  as 
the  "adiabatic  calorimeter"  has  been  used  for  several  years  in  the 
laboratory  of  T.  W.  Richards,  of  Harvard  University. 

The  two  chief  laws  of  thermochemistry  are: 

The  First  Law,  advanced  by  Laplace  and  Lavoisier  before  the  first 
law  of  energy  had  been  formulated,  states  that: 

As  much  heat  is  given  out  in  the  formation  of  a  substance  as  is  required 
to  separate  it  into  its  constituents. 

The  Second  Law  (Law  of  Hess)— 

No  matter  how  many  stages  there  are  in  a  given  chemical  reaction,  the 
quantity  of  heat  liberated  depends  upon  the  initial  and  final  states. 

This  is  sometimes  referred  to  as  the  "Constancy  of  the  Heat  Sum." 
Berthelot  proposed  a  Third  Law — 

"Reactions  go  in  the  direction  in  which  there  will  be  the  greatest 
evolution  of  heat." 

Whereas  this  has  considerable  value  as  a  sort  of  working  "rule,"  it  is 
not  to  be  considered  as  a  law,  for  it  is  now  known  that  "free"  energy,  and 
not  "  total "  energy,  determines  the  direction  of  the  reaction. 

*NoTE. — The  "  heat  of  reaction  "  is  equal  to  the  difference  of  the  sum  of  the  heats 
of  formation  of  the  original  substances  and  the  sum  of  the  heats  of  formation  of  the 
final  substances— the  heat  of  formation  of  the  elements  being  counted  zero. 


CHAPTER  V. 

SCIENCE— ITS  METHODS  OF  DEVELOPMENT  ; 
CLASSIFICATION. 

DEVELOPMENT    OF    SCIENCE. 

Men  have  been  led  naturally  to  investigate  the  causes  of  the  various 
phenomena  observed  in  the  physical  universe.  As  the  result  of  these 
investigations  we  have  our  several  sciences. 

The  initial  step  in  the  development  of  any  science  is  the  securing  of 
knowledge  by  observation  and  experiment.  By  an  experiment  is  usually 
meant  the  process  of  exhibiting  certain  phenomena  under  conditions 
proposed  and  controlled  by  the  experimenter.  In  the  conduct  of  these 
experiments  and  in  all  investigation,  we  proceed  upon  the  truth  of  wh  at 
is  now  known  as  a  general  maxim  in  physical  science — "the  constancy  of 
nature,"  i.e.,  "like  causes  produce  like  effects"  and  that,  irrespective  of 
time  and  place.  The  facts  secured  by  such  procedure,  isolated,  and 
apparently  unrelated,  become  the  data  of  science.  A  fact  is  that  which 
has  been  demonstrated  with  such  accuracy  as  to  leave  no  reasonable 
doubt  of  its  truthfulness.  As  the  investigation  of  nature  proceeded, 
the  multitude  of  facts  became  cumbersome  and  unwieldy.  Men  began 
to  sift  this  diverse  and  apparently  unrelated  data,  and  found  much  that  was 
common  to  many  phenomena.  In  the  stating  of  these  universal  and  valid 
relationships  they  formulated  what  are  now  known  as  laws.  After  ob- 
serving the  gravitational  tendency  of  many  bodies,  which  observations 
in  themselves  constituted  facts,  the  universal  relationship  was  recognized 
and  enunciated  in  the  "Universal  Law  of  Gravitation."  A  law  is  a 
statement  of  a  constant  relation  common  to  a  number  of  phenomena;  it 
is  the  expression  of  a  relation  among  facts.  "  When  two  substances  agree 
in  some  few  properties  they  also  agree  with  regard  to  all  other  properties." 
— Ostwald.  This  statement  of  the  relationship,  as  a  summary  of  ob- 
served facts,  is  called  a  law.  It  makes  no  endeavor  to  explain.  The 
student  should  understand  that  "it  does  not  govern  or  state  what  will 
happen;  it  predicates  an  invariable  relation." — Nernst.  There  are  no 
exceptions  to  laws;  if  there  is  an  exception  to  a  law  it  is  no  law.  The 
word  "rule"  has  been  frequently  confused  with  the  word  "law"  in  this 
latter  sense.  A  "rule"  is  merely  a  statement  of  the  usual  order  of  events 
— the  customary  procedure  and  condition  of  things,  as,  the  earth  turns 
upon  its  axis  once  a  day.  "Exceptions  to  rules  do  not  destroy  them  as 
rules,  in  fact  we  say  'the  exception  proves  the  rule.'" — Richards,  T.  W. 

It  is  interesting  to  note  the  two  methods  by  which  these  relationships 


SCIENCE — ITS    METHODS    OF    DEVELOPMENT;    CLASSIFICATION.          41 

have  been  discovered.  When  the  individual,  by  experiment  and  obser- 
vation, gathers  data  concerning  phenomena  among  which  he  suspects  a 
relationship,  and  arranges  it  empirically  and  then  endeavors  by  experi- 
ment to  verify  his  suspicions,  he  is  said  to  resort  to  the  inductive  method, 
sometimes  called  the  empirical  method. 

The  other  method  used  in  an  endeavor  to  discover  relationships  is 
known  as  the  deductive  or  theoretical  method,  and  is  largely  a  mental  proc- 
ess. It  consists  in  starting  with  merely  a  conception  regarding  the  re- 
lations among  phenomena,  then  by  pure  speculation  to  make  new  de- 
ductions, which  must  be  verified  by  research. 

The  inductive  method  leads  us  to  definite  and  usually  reliable  results. 

The  deductive  method  is  more  fascinating  and  enables  us  to  penetrate 
more  deeply  into  the  relationships,  but  we  must  be  very  careful  and 
cautious  in  the  choice  of  the  original  concept  which  many  times  can  not  be 
verified  by  a  direct  test. 

Conceptions  of  this  kind  which  can  not  be  directly  verified  are  called 
hypotheses.  The  imaginative  element  predominates  in  these  concepts 
which  are  mere  assumptions  of  "the  existence  of  conditions  of  which 
we  can  have  no  direct  experimental  evidence."  An  hypothesis  is  an 
effort  to  explain  the  mechanism,  the  "why"  of  certain  phenomena,  and 
to  correlate  same,  that  from  this  general  proposition  new  conclusions  may 
be  drawn.  This  new  body  of  conclusions  constitutes  a  theory,  which 
should  be  tested  by  experimental  methods.  A  prediction  is  merely  a 
guess  and  can  be  tested. 

In  this  connection  our  attention  is  called  to  the  scientific  and  systematic 
methods — the  two  being  commonly  confused,  although  they  are  entirely 
distinct.  Science  consists  in  discovering  the  laws  to  which  phenomena 
conform  and  is  exemplified  by  such  studies  as  chemistry,  biology,  mathe- 
matics and  similar  analytical  studies.  System  consists  merely  in  the 
classification  of  phenomena  as  illustrated  by  grammar,  history,  etc. 
Natural  history  preceded  biology,  the  former  merely  classified;  the  latter 
has  become  a  true  science  by  the  introduction  of  analytical  ideas  of 
the  relations  of  heredity,  environment,  etc.  Many  systematic  studies 
have  now  given  away  to  the  truly  scientific. 

CLASSIFICATION    OF    SCIENCE. 

In  the  course  of  time  knowledge  became  so  extensive  and  diverse  that 
it  was  evident  that  the  great  bulk  of  related  knowledge  which  might 
properly  be  called  Natural  Science  must  be  divided  for  convenience  into 
a  number  of  more  limited  ones.  These  various  divisions  of  the  one  large 
study  or  science  are  called  the  "natural  sciences."  The  distinction 
betwe.en  "living  matter"  and  "non-living  matter"  is  made  the  basis  for 
a  division  of  these  natural  sciences  into  two  great  groups  known  as  the 
Biological  Sciences  and  the  Physical  Sciences.  The  Biological  Sciences 
deal  with  matter  in  the  living  condition,  while  the  Physical  Sciences  has 
to  do  primarily  with  matter  in  the  lifeless  condition. 


42  EXPERIMENTAL  CHEMISTRY. 

Matter. 

I.     Organic — Biological  Sciences. 
A.     Biology  (general), 
a  i.     Botany, 
a  2.     Zoology. 

II.     Inorganic — Physical  Sciences. 

A.     Physics  and  Chemistry. 

a  i.     Astronomy  (Physics  of  the  heavens), 
a  2.     Meteorology, 
a  3.     Geology,  etc. 

As  will  be  observed  from  the  above  outline  it  is  possible  to  extend 
our  original  classification  into  what  are  known  as  the  General  or  Specula- 
tive Sciences  and  the  Special  or  Descriptive  Sciences.  Physics  and 
Chemistry  occupying  themselves  \vith  the  study  of  general  properties  and 
"transformations"  of  bodies  regardless  of  external  form,  and  dealing 
with  the  substance  only,  are  called  the  Speculative  or  General  Sciences. 
Geology,  Astronomy,  Meteorology,  etc.,  considering  distinct  classes  of 
bodies  in  reference  to  form,  classification  and  distinguishing  characteristics, 
are  known  as  the  Descriptive  or  Special  Sciences. 

Chemistry  is  a  descriptive  science  in  so  far  as  it  considers  the  external 
properties  of  chemical  substances. 

Again  all  knowledge  may  be  gathered  under  the  respective  headings  of 
the  Abstract  Sciences  and  the  Concrete  Sciences. 

The  Abstract  Sciences,  including  logic,  mathematics,  etc.,  are  not  con- 
cerned primarily  with  matter,  but  with  abstract  conceptions.  The 
Concrete  Sciences  are  concerned  with  bodies,  i.e.,  "definite  aggregation  of 
matter,"  whether  living  or  lifeless.  Zoology,  Astronomy,  etc.,  would  be 
included  under  this  head.  Physics  and  Chemistry,  having  to  do  with 
both  the  abstract  and  the  concrete,  are  frequently  named  the  "  abstract- 


CHAPTER  VI. 
QUANTITATIVE  RELATIONSHIPS. 

LAWS  AND  CHEMICAL  EQUIVALENTS. 

As  a  science  progresses  in  accuracy  it  progresses  in  the  true  scientific 
spirit.  We  have  been  led  to  the  conclusion  in  the  foregoing  work  that 
definite  quantities  of  any  substance  under  defined  conditions  possess  a 
definite  quantity  of  internal  energy.  The  question  might  rise  logically, 
as  to  whether  the  "factors"  and  "products,"  as  referred  to  matter,  in- 
volved in  a  chemical  change  bear  any  such  definite  relation  to  one  another, 
that  is,  when  substances  combine  to  form  new  compounds,  (a)  does  the 
matter  in  the  system  sustain  an  appreciable  loss  or  gain,  (b)  do  the  re- 
acting substances  combine  in  any  definite  ratio  by  weight  or  volume,  (c) 
do  the  constituents  of  the  new  substance  bear  to  one  another  a  fixed  ratio 
by  weight  or  volume,  (d)  does  variation  in  the  conditions  cause  a  variation 
in  any  possible  ratio  in  which  substances  may  combine  with  one  another? 
To  aid  us  in  understanding  the  conclusions  which  have  been  reached  in 
regard  to  these  questions,  as  the  result  of  much  patient  and  painstaking 
labor,  let  us  perform  a  series  of  experiments. 

Experiments  are  of  two  kinds,  qualitative  and  quantitative.  The  former 
are  resorted  to  in  order  to  illustrate  a  property  or  group  of  properties. 
We  consider  only  what  takes  place.  Quantitative  experiments,  for  ex- 
ample, accept  the  fact  that  gases  are  compressible,  then  endeavor  to 
determine  the  exact  amount  of  compressibility  for  each  increment  of 
pressure.  They  seek  to  determine  any  possible  mathematical  relation 
of  cause  and  effect. 

Experiment  I. — Conservation  of  Mass. 

(a)  (L.  T.) — Analysis  of  Water.  Counterpoise  an  electrolysis  apparatus 
upon  a  scale  pan,  or  better,  fit  a  rubber  stopper  containing  three  holes 
into  a  250  cm. 3  flask;  half  fill  the  flask  with  water  slightly  acidulated 
with  sulphuric  acid;  insert  two  electrodes  through  holes  in  cork,  reserving 
third  hole  for  a  glass  tube  with  rubber  tubing  and  pinch  cock  attached ; 
by  means  of  this  latter  and  an  air-pump  exhaust  air  above  water  from 
flask ;  close  pinch  cock ;  invert  flask ;  mark  height  of  contents ;  counterpoise 
apparatus  on  balance ;  turn  on  the  electric  current ;  note  bubbles  of  gas 
arising  from  electrodes.  Caution:  Don't  continue  process  of  decompo- 
sition too  long.  Counterpoise  apparatus  again ;  any  gain  or  loss  in  weight  ? 
Is  the  sum  of  the  weights  of  the  "products"  equal  to  the  sum  of  the 
weights  of  the  "factors?" 

Water  — »  Hydrogen  +  Oxygen. 

43 


44 


EXPERIMENTAL  CHEMISTRY. 


(b)  (L.  T.) — Synthesis  of  Water.  Perform  this  operation  by  means  of  a 
Hoffman  eudiometer  (Figs.  5  and  6).  Introduce  into  apparatus  about  equal 
volumes  of  hydrogen  and  oxygen ;  observe  volume  of  the  mixture  of  gases. 

Note. — It  is  well  to  use  an  excess  of  one  or 
the  other  of  the  gases  so  that  when  the  con- 
traction of  volume  of  gases  occurs,  there 
will  be  a  cushion  of  gas  between  top  of  glass 
tube  and  the  mercury.  What  deductions,  as 
regards  the  relation  of  products 
and  factors,  can  you  make  ? 

Hydrogen  +  Oxygen  —  Water. 

As  the  result  of  much  refined 
work  it  has  been  concluded  that 
regardless  of  chemical  change 
there  is  "conservation  of  mass" 
in  a  limited  system.  It  is  quite 
possible  that  Van  Helmont 
(1577-1644)  was  acquainted 
with  the  principle.  The  law 
was  first  enunciated  by  Lavoi- 
sier (1789).  It  is  commonly 
spoken  of  as  the  law  of  the  "in- 
destructibility of  matter." 

Experiment      II. — Combina- 
FIG.  5.  tion  in  Definite  Proportions  by 

Weight. 

(a)  Synthesis  of  copper  oxide.  Clean  and  dry  a  small 
porcelain  crucible;  wreigh  it;  put  into  it  about  i  gram  of 
powdered  copper;  secure  the  exact  weight;  place  crucible  on 
pipe -stem  triangle  which  may  be  supported  by  tripod  or  ring 
clamp.  Heat  crucible  and  contents  to  redness  in  a  Bunsen 
flame  for  fifteen  minutes ;  remove  flame  and  let  crucible  cool ; 
weigh;  make  record  of  weight.  Repeat  operations  until 
weight  becomes  constant. 

Copper  +  Oxygen  — *  Copper  Oxide. 
Tabulate  your  data  as  follows: — 

Weight  of  crucible  and  copper o. ooo 

Weight  of  crucible o .  ooo 

FIG  6. 


Weight  of  copper o .  ooo 

Weight  of  crucible  and  copper  before  heating 

Weight  of  crucible  and  copper,  after  heating 

Weight  of  oxygen,  combined  with  copper 


Eudiometer. 


o.ooo 
o.ooo 


o.ooo 


QUANTITATIVE    RELATIONSHIPS.  45 

Calculate  how  much  oxygen  would  combine  with  a  gram  of  copper. 
What  is  the  ratio  of  the  combining  weights  of  copper  and  oxygen  under 
above  conditions  ?  What  is  the  class  average  ? 

(b)  Optional.     Synthesis  of  magnesium  oxide.     Prepare  crucible  and 
cover  as  above;  wreigh  into  it  about  300  mg.  of  magnesium  ribbon;  use  for- 
ceps and  scissors  for  manipulation  of  ribbon;  make  a  record  of  the  exact 
weights  as  above.     Place  covered  crucible  on  triangle  and  cautiously 
heat  it  to  a  dull  redness — then  quickly  removing  flame,  raise  cover  of 
crucible  and  observe  whether  the  magnesium  starts  to  glow,  which  is 
desired, — if  it  does,  quickly  replace  cover  so  that  none  of  the  magnesium 
oxide  escapes.     This  operation  of  heating  to  redness,  removing  flame 
and  lifting  cover  must  be  repeated  until  a  constant  weight  is  secured 
and  the  contents  of  crucible  have  a  white  or  grayish  color.     Tabulate  data 
and  calculate  ratio  of  combining  weights  as  in  (a). 

Magnesium  +  Oxygen  — *  Magnesium  Oxide. 

(c)  Optional.     Synthesis    of   iron    oxide.     Pure   iron   wire   dissolved 
in  nitric   acid  gives  a  compound  which  on  being  heated  sufficiently,  is 
converted  into  iron  oxide.     Into  a  thoroughly  cleaned  and  dried  evap- 
orating dish  which  has  been  weighed,  introduce  about  750  mg.  of  iron 
wire.     (Instructions.)     Record  exact  weight  of  wire  used.     Place  cover 
glass  on  dish;  add  10  cm.3  of  dilute  nitric  acid.     Now  fill  a  beaker  nearly 
half  full  of  water  and  heat  the  water  until  it  boils  vigorously ;  place  covered 
evaporating  dish  on  beaker  so  that  the  latter  serves  as  a  steam  bath ;  keep 
cover  on  until  all  of  the  iron  wire  has  dissolved,  then  remove  cover  and 
evaporate  to  dryness;  the  direct  flame  may  now  be  applied  to  the  dish; 
continue  its  use  until  a  constant  weight  is  secured.     As  long  as  red  fumes 
are  observed  rising  from  the  dish,  it  will  be  unnecessary  to  cool  and 
weigh.     The  powder  remaining  in  dish  is  iron  oxide.     Tabulate  data  and 
calculate  as  in  (a).     What  is  the  ratio  of  the  combining  weights  ? 

It  is  quite  obvious  in  these  experiments  that  we  have  not  been  able  to 
get  substances  to  combine  in  any  promiscuous  ratio;  on  the  contrary  the 
combining  substances  have  united  in  certain  fixed  ratios  by  weight.  A 
generalization  of  this  principle  is  known  as  the  "Law  of  Definite  Pro- 
portions." It  was  practically  confirmed  and  announced  by  Proust  in 
.1806,  and  came  largely  as  the  result  of  a  controversy  between  Berthollet 
and  Proust.  The  law  as  frequently  stated  is, 

The  same  compound  always  contains  the  same  elements  combined  together 
in  the  same  proportion  by  weight. 

Experiment  III. — Constancy  of  Composition. 

Liberation  of  oxygen  by  decomposition  of  potassium  chlorate. 

Note. — A  small  quantity  of  potassium  chlorate  may  be  heated  in  a  test 
tube  and  the  evolved  gas  tested  for  oxygen. 

Clean  and  dry  and  heat  a  porcelain  crucible  and  lid ;  cool  in  a  desiccator; 
weigh  accurately  and  record  weight ;  introduce  into  the  crucible  an  accu- 


46  EXPERIMENTAL  CHEMISTRY. 

rately  weighed  quantity,  say  1.2  gram,  of  powdered  potassium  chlorate 
which  has  been  previously  dried  at  a  temperature  of  about  100°  C.  for  15 
to  20  min.  Place  crucible  and  contents  on  a  pipe-stem  triangle  and  heat 
gently  in  order  to  prevent  spattering ;  if  any  of  the  salt  is  deposited  upon  the 
cover,  the  latter  must  be  cooled,  and  the  salt  returned  to  crucible.  When 
the  mass  solidifies  and  action  has  apparently  stopped,  increase  the  heat 
until  a  perfectly  white,  non-crystalline  mass  which  is  not  altered  by  a 
further  increase  of  heat,  results.  Remove  cover  and  heat  strongly  for  a 
few  minutes;  cool  in  desiccator  and  weigh  as  before.  Record  weight. 
Heat  again  for  10-15  min.;  cool;  weigh.  Repeat  until  weight  is  practi- 
cally constant.  Tabulate  all  data  as  in  a  previous  experiment.  Calcu- 
late the  percentage  of  oxygen  in  potassium  chlorate. 

Above  result  indicates  in  the  case  of  this  compound,  potassium  chlorate, 
that  a  definite  proportion  by  weight  of  oxygen  can  be  secured  from  a  defi- 
nite weight  of  the  compound.  Further  experimentation  confirms  this 
principle,  known  as  the  "law  of  constancy  of  composition."  It  is  but 
a  special  statement  of  what  law  ? 

Experiment  IV. — Combination  of  Two  Substances  in  Varying  Propor- 
tions by  Weight. 

(a)  Optional.     Synthesis  of  mercuric  iodide  (Hg  I2),  as  compared  with 
synthesis  of  mercurous  iodide  (Hg  I).     (Instructions  from  the  assistant.) 

(b)  A  qualitative  experiment  designed  to  show  in  a  general  way  the 
difference  between  the  properties  of  two  compound  substances,  each  of 
which  is  composed  of  the  same  elementary  substances  but  in  different 
proportions  by  weight.     Synthesis  of  nitrogen  dioxide  (N2  O2)  as  com- 
pared with  the  synthesis  of  nitrogen  tetroxide  (N2  O4).     Assemble  parts 
of  a  gas  generating  flask  equipped  with  delivery  tube ;  place  about  10  grams 
of  copper  turnings  in  the  flask;  arrange  delivery  tube  so  that  it  dips  into 
a  vessel  of  water  (pneumatic  trough).     Pour  through  the  funnel  tube 
about  15  cm. 3  of  nitric  acid  diluted  with  its  own  volume  of  water,  then 
add  a  few  cm.3  of  concentrated  nitric  acid  at  a  time  until  there  is  a  rapid 
evolution  of  gas;  after  gas  has  been  evolved  for  a  few  minutes  collect 
several  bottles  of  the  colorless  gas  by  water  displacement;  place  a  glass 
plate  or  a  piece  of  wet  paper  over  mouth  of  bottle  on  lifting  it  from  water; 
remove  cover  so  that  oxygen  of  air  can  come  into  intimate  contact  with 
the  colorless  gas.     Does  the  gas  in  the  bottle  become  brownish-red  in 
color  ?     The  colorless  gas  is  a  combination  of  nitrogen  and  oxygen  in  the 
ratio  by  weight  of  14  to  16;  the  brownish-red  gas  is  composed  of  the  same 
gases  but  in  the  ratio  by  weight  of  14  to  32.     Were  the  conditions  under 
which  two  compounds  formed  identical  ?     Explain. 

This  is  an  illustration  of  the  "Law  of  Multiple  Proportions." 
//  two  elements  unite  in  more  than  one  proportion,  forming  two  or  more  com- 
pounds, the  different  weights  of  one  oj  the  elements,  which  in  the  different  com- 
pounds are  united  with  identical  amounts  oj  the  other,  bear  a  simple  ratio 
to  one  another. 


QUANTITATIVE    RELATIONSHIPS. 


47 


It  is  not  an  exception  to  the  law  of  definite  proportions  which  applies 
to  either  one  of  the  compounds  formed.  The  conditions  under  which 
each  compound  was  formed  were  different.  See  under  "  Copper."— Exp., 
Preparation  of  Cuprous  and  Cupric  Chlorides. 

Experiment  V.— Combining  Volumes  of  Gases. 

(a)  Synthesis  of  water  from  hydrogen  and  oxygen  gases.  Introduce  30 
cm.3  of  each  of  the  two  gases  into  a  Hoffman  eudiometer,  U-form  (Fig.  5). 
Observe  proper  precautions  to  collect  gases  at  atmospheric  pressure. 
Open  lower  stop  cock  permitting  some  of  the  mercury  to  run  out,  thus  al- 
lowing gases  to  expand;  pass  spark,  and  then  bring  mercury  columns 
to  same  level.  Read.  Repeat  using  25  cm.3;  2o  cm.3,  of  each  gas.  RC_ 
cord  all  data.  What  is  the  average  of  the  ratios  of  the  combining  volumes  ? 


FIG.  7. 

Note. — Before  passing  spark  through  mixture  of  gases  it  is  well  to  place 
thumb  on  open  tube  to  prevent  the  mercury  from  being  forced  out. 

(b)  Synthesis  of  steam  from  hydrogen  and  oxygen  gases. 

Introduce  25  cm.8  of  each  of  the  two  gases  into  an  eudiometer  (Fig.  7) 
which  is  inclosed  in  a  glass  tube  which  may  serve  as  a  steam  jacket ;  the  eudi- 
ometer should  be  connected'with  a  leveling  bulb  and  the  gases  collected  over 
mercury;  the  gases  must  be  perfectly  dry,  as  should  the  mercury;  after 
gases  have  been  collected,  pass  steam  through  jacket  until  entire  system 
is  at  the  temperature  of  steam ;  observe  volume  of  mixture  of  gases  when 
mercury  in  eudiometer  is  at  same  height  as  it  is  in  the  leveling  bulb; 
lower  leveling  bulb  permitting  gases  to  expand ;  clamp  tubing  below  eu- 
diometer; pass  spark  through  mixture  of  gases;  bring  mercury  columns 
to  the  same  height  again;  note  volume.  Deduct  one-fourth  the  volume 
of  the  mixture  of  gases  at  100°  C.  from  this  final  volume  (Why?).  What 
relation  does  this  volume  bear  to  three-fourths  of  the  volume  of  the  mix- 
ture of  gases  at  the  temperature  of  the  steam  ? 


48  EXPERIMENTAL  CHEMISTRY. 

Note. — If  more  desirable  to  teacher,  the  exact  combining  volumes 
may  be  used  as  prepared  by  electrolysis  of  water. 

Experiment  has  shown  that  two  volumes  of  hydrogen  gas  combine 
with  one  volume  of  oxygen  gas  to  produce  water.  Further  it  has  been 
shown  that  when  the  quantity  of  water  has  been  converted  into  steam, 
that  its  volume  is  to  the  sum  of  the  combining  volumes  of  hydrogen  and 
oxygen  at  the  same  temperature,  as  2  to  3. 

These  and  similar  experimental  results  have  given  rise  to  a  law,  known 
as  the  "Law  of  Gaseous  Volumes"  or  "The  Law  of  Gay-Lussac. " 

When  chemical  action  takes  place  between  gases,  either  elements  or  com- 
pounds, the  gaseous  product  bears  a  simple  relation  to  the  volume  of  the 
reacting  gases. 

A  relationship  which  may  be  noted  at  once  by  a  study  of  the  foregoing 
experiments  is  stated  as  the  Law  of  Reciprocal  Proportions,  or  Law  of 
Equivalent  Proportions. 

The  weights  of  different  elements  which  combine  separately  and  with  one 
and  the  same  weight  of  another  element,  are  either  the  same  as  or  are  simple 
multiples  of,  the  weights  oj  these  different  elements  which  combine  with 
each  other. 

It  might  be  called  the  "law of  inter-equivalence  of  equivalent  weights." 
An  application  of  this  principle  may  be  made  as  follows: — 63.6  parts  by 
Aveight  of  copper  or  24.3  parts  by  weight  of  magnesium  combine  with  1 6  parts 
by  weight  of  oxygen;  2  parts  by  weight  of  hydrogen  also  combine  with  1 6 
parts  by  weight  of  oxygen ;  therefore,  it  is  obvious  that  63.6  parts  of  copper, 
24.3  parts  of  magnesium  and  2  parts  of  hydrogen  are  chemically  equiva- 
lent, that  is,  these  proportions  by  weight  satisfy  the  chemical  affinity  of  one 
another.  Again  it  is  to  be  observed  that  i  part  by  weight  of  hydrogen  is 
chemically  equivalent  to  31.8  parts  by  weight  of  copper,  12.2  of  magne- 
sium, and  8  of  oxygen.  Inasmuch  as  hydrogen,  the  lightest  of  all  ele- 
ments, has  the  smallest  combining  weight,  Dalton  proposed  to  make 
this  weight  the  standard,  and  express  all  combining  weights  in  terms  of  it. 
Berzelius  proposed  oxygen,  which  is  about  16  times  as  heavy.  These 
weights  of  elements  which  are  chemically  equivalent  to  i  part  by  weight 
of  hydrogen  are  called  chemical  equivalents.  More  accurately,  that  weight 
of  an  element  which  combines  with  or  displaces  8  parts  of  oxygen  or  1.0075 
parts  of  hydrogen  by  weight  is  called  the  equivalent  weight.  Specifically, 
it  is  the  number  of  grams  oj  an  element  which  combines  with  or  replaces 
1.0075  grams  of  hydrogen  or  8  grams  oj  oxygen. 

IF  the  student  finds  this  word  "equivalent"  difficult  to  understand  in 
this  connection,  some  satisfaction  may  be  secured  by  knowing  that  this 
same  word  delayed  the  development  of  chemical  theory  many  years. 

Experiment  VI.  —Equivalent  Weight  of  Zinc  by  Displacing  of  Hydrogen. 

See  Exp.  under  "Hydrogen." 


CHAPTER  VII. 
OUTLINES. 

THEORIES,  FORMULAE,   VALENCE,  EQUATIONS  AND  UNITS. 

Structure  of  Matter. j — Continuous  or  discontinuous?  Matter  bears 
evidence  of  possessing  a  granular  structure. 

Statement  and  Historical  Development  of  the  Kinetic-Molecular 
Hypothesis. 

Definitions  of  a  Molecule. 

(1)  A  molecule  is  the  smallest  particle  of  matter  which  can 
exist  alone  and  retain  the  properties  of  a  substance. 

(2)  "  Molecules  are  the  imaginary  units  of  which  bodies  are 
the  aggregates." — Smith. 

(3)  A  molecule  is  the  physicist's  unit.* 

Application  of  Kinetic-Molecular  Hypothesis  to, 

(1)  Solids, 

(2)  Liquids, 

(3)  Gases. 

Boyle's  Law. 

Charles'  Law  (Dalton — Gay-Lussac — Charles). 

Avogadro's  Rule. 

Gas  Law  (PV  =  RT). 

Critical  Phenomena.— Experiments  with  " critical"  tubes. 

(a)  Critical  Temperature  (H,— 238°C.),  (O,— n8°C.),  (Ether, 

.i95°C.). 

(b)  Critical  Pressure  (H,  20  atmos.),   (O,  51  atmos.),  (Ether, 

36  atmos.). 

(c)  Critical  Point. 

(d)  Critical  Volume. 

Combining  Volumes.— Recall  Exp.  and  Law  as  discussed  under 
"Quantitative  Relationships."  Two  volumes  of  hydrogen,  say  40  cm. 3 
combine  with  one  volume,  20  cm.3  of  oxygen  to  form  two  volumes,  40  cm.3 
of  steam  (water). 

n  n         +         n  n  n 

2  Vol.  Hydrogen,     i  Vol.  Oyxgen.     2  Vol.  Steam, 
(i  molecule  of  oxygen  must  divide.) 

n      +       n      —      n  n 

Hydrogen.     Chlorine.     Hydrogen  Chloride. 

(Both  molecules  must  divide.) 

*  The  modern  physicist  will  probably  insist  that  the  "  electron  "  is  the  unit. 
4  AO 


$0  EXPERIMENTAL  CHEMISTRY. 

Statement  and  Historical  Development  of  Atomic  Theory. 

Definitions  of  an  Atom. 

(1)  An   Atom  is   the   smallest  particle   of   matter  which   can 
participate  in  a  chemical  reaction. 

(2)  Atoms  are  the  imaginary  units  of  which  molecules  are  the 
aggregate. 

(3)  An  atom  is  the  chemist's  unit. 

(4)  An  atom  always  represents  a  distinct  variety  of  matter. 

Statement  of  J.  J.  Thomson's  Electronic  Theory  of  Matter.  See 
Thomson's  "  Electricity  and  Matter." 

Application  of  Atomic  Theory  to, 

(1)  Chemical  Reaction. 

(2)  Quantitative  Relationships. 

Atomic  Weight  "  is  the  number  which  represents  the  smallest  mass  oj  an 
element  which  is  known  to  participate  in  a  chemical  change,  relative  to  the 
smallest  weight  oj  hydrogen  which  can  so  junction" — Newth;  or,  "it  is 
the  unit  oj  weight  actually  used  in  expressing  the  proportions  oj  each  ele- 
ment in  all  oj  its  compounds" — Smith. 

Note. — Hydrogen  has  been  practically  abandoned  in  favor  of  oxygen 
as  the  standard  of  atomic  weights. 

Nomenclature  of  Atomic  Theory. 
Symbols  (Hg,  O,  H,  etc.). 
Molecular  Formulae. 

(a)  Simple  Molecules  (H2,  O2,  O3,  Hg,  As4,  S8). 
(6)  Compound  Molecules  (FeS,  HgO,  H2O,  HC1,  MgO). 

Molecular  Weight  is  equal  to  the  sum  oj  the  atomic  weights  oj  all  oj  the 
atoms  oj  which  the  molecule  is  composed. 

DETERMINATION   OF   MOLECULAR    FORMULA    AND    EQUATIONS. 
Equations. — 

(a)  Matter  or  Mass  Equations. 

EXAMPLES: 

H2  +  O  -*  H2O. 
H2  +  C12  —  2HC1. 
Mg  +  O  — »  MgO. 

(b)  Energy  Equations. 

EXAMPLES: 

H2   +  O    —  H2O    +  68,000  cal. 

H2   +  C12  —  2HC1  +  44,000  cal. 

C     +  O2  —  CO2    +  97,000  cal. 

H2   +  I2    — *  2HI    —  12,000  cal. 

Hg  +  O    — »  HgO   +  22,000  cal. 


OUTLINES.  51 

When  a  body  or  substance,  by  virtue  of  its  position  or  its  motion  or 
its  state,  is  capable  of  doing  work  it  is  said  to  possess  energy ;  if  resistance 
is  overcome,  work  is  done.  Atoms  do  not  exist  as  individuals  except 
in  a  few  instances  (Hg,  Na,  K,  Zn,  Cd,  He,  A,  Kr,  Ne,  X);  they  are 
united  to  form  molecules,  simple  or  compound,  of  varying  degrees  of 
stability.  The  stability  of  this  equilibrium  is  evidently  determined  by 
some  force  operative  between  the  individual  atoms ;  for  want  of  a  better 
term  we  designate  this  force  "chemism"  or  " chemical  affinity"  (see 
Note  on  Energetics  of  Chemistry).  To  return  to  our  original  line  of 
thought,  it  is  evident  that  atoms  possess  energy,  because  in  uniting  they 
are  capable  of  performing  work.  Atoms  are  said  to  possess  "chemical 
energy"  (see  Note  on  Energetics  of  Chemistry)  in  virtue  of  their  tendency 
to  combine  with  other  atoms. 

It  is  probable  that  the  energy  liberated  when  atoms  combine  is  due  to 
a  transformation  of  all  or  a  part  of  this  chemical  energy  into  jree  energy, 
i.e.,  energy  which  is  available  for  the  performance  of  work.  If  certain 
kinds  of  atoms  have  a  great  chemical  affinity  for  one  another  they,  that 
is  the  system  which  they  represent,  possesses  a  large  amount  of  chemical 
energy  which  will  probably  be  converted  into  a  correspondingly  large 
amount  of  free  energy.  Molecules  formed  under  such  conditions  possess 
much  smaller  quantities  of  energy  than  their  constituent  atoms  originally 
possessed,  consequently  they  are  stable.  Again,  two  or  more  atoms 
may  possess  so  very  little  chemical  affinity  for  each  other  that  it  will  be 
necessary  to  contribute  energy  to  secure  their  combination;  molecules 
formed  under  these  conditions  will  have  a  tendency  to  convert  this 
acquired  energy  into  free  energy  and  therefore,  will  be  unstable. 

The  atoms  of  the  elements  show  great  variations  in  respect  to  the 
chemical  affinity  which  they  manifest  toward  one  another;  the  same  is 
true  as  regards  the  interaction  of  groups  of  atoms. 

Valence,  or  Quantivalence  is  the  atom  -fixing  or  replacing  power  of  an 
atom  of  an  element  in  terms  of  the  hydrogen  atom  which  is  considered  as 
having  a  valence  of  one.  Those  atoms  which  have  a  valence  of  one, 
two,  three,  etc.,  are  said  to  be  respectively  univalent,  bivalent,  trivalent,  etc. 

Multiple  Valence. — The  valence  of  an  elementary  atom  "is  not  an 
invariable  property  oj  the  atonij  but  each  of  the  observed  differences  as  to  its 
valence  is  an  invariable  property  of  some  particular  condition  oj  the  atom 
dependent  on  the  circumstances  in  which  it  is  placed." 


Formulae. 


Empirical,  Ex.,  H2O,  C2H4O2. 

Rational,  Ex.,  HC2H3O2. 

Dualistic,  Ex.,  H2O.SO3,  Na2O.SO3, 

H \          H\         H] 
Typical,    Cl  /     '      H  /  U'    H     N. 

H 


52  EXPERIMENTAL  CHEMISTRY. 

H  H  H-0          O 

/     /  \  / 

Structural  or  Graphic,  O      ,     N-H,     Na  — OH,  S 

\  \  /  \ 

H  H  H-O          O 

Space   (formulae   of   more   than  one  dimension,  to  represent  the  con- 
figuration of  a  molecule). 

Units  of  Measurement. 

Length — Cm.,  mm. 
Volume — Cm.3,  liter. 

Weight. 

(a)  Gram. 

(b)  Gram-molecule  (mole). 

(c)  Formula- Weight. 

Temperature — Centigrade  Scale. 

(a)  Boiling  Point  (ioo°C.). 

(b)  Freezing  Point  (o°  C.). 

(c)  Absolute  Zero  (—  273  °C.). 

Heat  Units— Calorie. 

(a)  Small  Calorie. 

(b)  Large  Calorie. 

(c)  Ostwald  Calorie. 

Calculations  in  Chemistry  (Stoichiometry). — Problems. 

Mg  +  O  —  MgO. 

24        1 6      40. 
KC1O3-*KC1  +  30. 
122.5         74-5       48. 

DETERMINATION    OF    MOLECULAR    AND    ATOMIC    WEIGHTS. 

Methods. 

(1)  Volumetric. — (Molar-volume  22.4  1.). 

(2)  Chemical. 

(3)  Specific  Heats — Atomic  Heat. 

Dulong  and  Petit's  Law — Neumann's  Rule. 

(4)  Isomorphism. 

(5)  Elevation  of  Boiling  Point. 

(6)  Depression  of  Freezing  Point. 

References. 

Atomic  Theory. — Any  large  text-book  in  chemistry.  Theoretical 
Chemistry. — Remsen.  Heroes  of  Science,  Chemists. — Muir.  Atomic 
Theory.— Wurtz.  Alembic  Club  Reprints,  No.  II. 

Avogadro's  Rule. — Chemical  Theory. — Dobbin  and  Walker.  Out- 
lines of  General  Chemistry. — Ostwald. 


OUTLINES.  53 

Critical  Phenomena. — Introduction  to  Physical  Chemistry. — Walker. 
Any  text-book  on  physical  chemistry. 

Valence. — Theoretical  Chemistry. — Remsen.  Theoretical  Chemistry. 
— Nernst. 

Specific  Heats  (Dulong  and  Petit's  Law). — History  of  Chemistry.— 
Ladenberg.  Outlines  of  Theoretical  Chemistry. — Meyer. 

Elevation  of  Boiling  Point,  etc. — Physical  Chemistry. — Walker. 
Physical  Chemistry. — Jones.  Theoretical  Chemistry. — Nernst.  Mod- 
ern Theory  of  Solution. — Jones.  See  later  work  under  the  subject 
of  "  Solutions." 

PROBLEMS. 

Note. — See  At.  Wt.  tables  in  Appendix. 

i.— What  is  the  molecular  weight  of  HgO,  H2O,  FeS,  MgO,  HC1, 
H2SO4,  HNO3,  NH4OH,  NaOH,  KOH? 

2. — Mg  -f  O  — »  MgO.  5  grams  of  magnesium  will  combine  with 
how  many  grams  of  oxygen  ?  Now  many  grams  of  magnesium  oxide  will 
be  formed? 

3. — How  many  grams  of  hydrogen  in  10  grams  of  H2O  ? 

4. — How  many  grams  of  oxygen  can  be  prepared  from  10  grams  of 
mercuric  oxide?  HgO  — >  Hg  +  O. 

5. — Fe  +  S  — »  FeS.  If  10  grams  each  of  sulphur  and  iron  are  placed 
in  the  test  tube  and  chemical  action  results  with  a  transfer  of  all  of  the 
iron  into  iron  sulphide,  how  much  sulphur  remains  in  the  uncombined 
condition  ? 

6. — How  much  oxygen  can  be  prepared  from  two  gram-molecules  of 
potassium  chlorate?  KC1O3— KC1  +  30. 

7. — A  certain  weight  of  KC1O3,  was  heated  until  completely  decom- 
posed as  per  above  equation.  The  residue  weighed  10.123  gm.  What 
was  the  original  weight  of  KC1O3,  and  how  much  oxygen  was  evolved? 
How  many  liters  of  oxygen  were  liberated  if  the  standard  weight  of  a 
1.  of  oxygen  is  1.4  gm.  ? 

8. — At  760  mm.  and  o°  C.  how  many  liters  of  gas  in  2  gm.  (molecular 
wt.)  of  hydrogen?  In  16  gm.  (.5  mol.  wt.)  of  oxygen?  In  32  gm. 
(mol.  wt.)  of  oxygen?  What  is  the  molar  volume? 

9. — (a)  5.6  1.  of  a  gaseous  substance  weights  4.5  gm.  approximately, 
at  standard  conditions ;  what  is  its  molecular  weight  ? 

(b)  If  |  by  weight  of  this  substance  is  hydrogen  and  f  oxygen, 
what  is  its  formula? 

10. — (a)  If  the  specific  heat  of  Ag  (silver)  is  approximately  .057, 
what  is  its  atomic  weight  ?  Of  Fe,  if  sp.ht.  is  .112  ? 

(b)  If  the  sp.ht.  of  Ag  Cl  (silver  chloride)  is  .089,  prove  that  the 
atomic  weight  of  Cl  is  nearer  35  than  70.  What  is  the  atomic  heat  of 
chlorine  ? 

n.— Write  the  structural  formula  for  NH3,  H2O,  MgO,  FeS,  HC1,  and 
H2SO4,  if  in  this  latter  compound  (sulphuric  acid)  sulphur  has  a  valency 
of  6. 


CHAPTER  VIII. 
OXYGEN. 

Symbol— O.     At.  Wt.  16  (15.9). 

Oxygen  is  an  elementary  substance  belonging  to  the  class  previously 
designated  as  "  non-metals."  It  was  first  discovered  by  Priestly,  who 
in  1774  prepared  it  "by  heating  'red  precipitate'  (red  oxide  of  mercury) 
in  the  focus  of  a  burning  glass  exposed  to  the  sun's  rays."  It  is  quite 
probable  that  Scheele,  a  Swedish  chemist,  had  previously  obtained  it, 
but  he  delayed  publication  of  his  results  until  1777.  Lavoisier,  a  French 
chemist,  named  the  new  element  "oxygen"  which  signifies  acid-pro- 
ducer. He  believed  that  all  acids  owe  their  characteristic  properties  to 
the  presence  of  oxygen.  Cavendish  pointed  out  that  many  substances 
which  do  not  contain  oxygen  possess  acid  properties.  We  now  know 
that  Lavoisier's  view  was  incorrect. 

Occurrence. — Oxygen  is  the  most  abundant  element  upon  our  planet; 
it  forms  about  47  per  cent,  of  the  solid  portion  of  the  earth  and  consti- 
tutes 20  per  cent,  of  the  earth's  atmosphere.  Including  its  occurrence 
in  the  ocean,  vegetable  and  animal  forms,  etc.,  it  constitutes  50  per 
cent,  of  the  total  substance  of  the  globe.  It  occurs  most  abundantly 
in  the  free  state  in  the  air. 

Physical  Properties. — .Oxygen  is  a  colorless,  odorless  and  tasteless 
gas;  it  is  slightly  heavier  than  the  air;  its  density  as  compared  with  air 
which  is  considered  as  i  (physical  standard),  is  1.105.  One  liter  of 
oxygen  at  standard  conditions  weighs  1.429  gm.  It  is  slightly  soluble  in 
water,  100  volumes  dissolving  4.  i  volumes  of  oxygen  at  5  °  C.  Its  critical 
temperature  is  — 118°  C. ;  its  critical  pressure,  50  atmospheres.  Liquid 
oxygen  boils  at  — 182°  C.  at  740  mm.  pressure  (Dewar  and  Fleming); 
it  has  a  pale-blue  color;  its  sp.  gr.  at  182.5°  C.  is  1.13. 

Chemical  Properties. — By  chemical  properties  is  meant  "reaction 
properties"  Oxygen  possesses  an  almost  universal  affinity,  i.e.,  "it 
forms  chemical  compounds,  called  oxides,  with  all  other  elements  ex- 
cepting fluorine  and  bromine  and  it  will  combine  with  the  latter  element 
provided  some  metal  is  also  a  constituent  of  the  compound." — Freer. 
Much  heat  is  liberated  during  many  of  these  oxidations,  i.e.,  formation  of 
oxides.  This  indicates  that  oxygen  possesses  considerable  chemical 
energy.  Substances  which  undergo  oxidation  also  possess  a  certain 
amount  of  chemical  energy  in  the  presence  of  oxygen,  and  when  they 
combine  with  the  latter  element  this  energy  is  converted  into  heat. 
Oxygen  is  ordinarily  bivalent. 

54 


OXYGEN.  55 

Experiment  I. — Sources  of  Oxygen. 

(a)  Place  a  small  quantity  of  mercuric  oxide  (HgO)  in  a  small  test 
tube.     Note  its  color.     Heat  bottom  of  tube  strongly  and  while  doing 
so,  introduce  a  glowing  splinter  of  wood.     Results?     Continue  to  heat 
until  sides  of  t.t.  show  a  metallic  lustre.     Explain.     Write  the  equation 
for  this  reaction.     When  the  test  tube  is  set  aside  and  permitted  to  cool, 
does  the  powder  in  the  bottom  of  the  tube  regain  its  original  color? 
Explain.     Write  the  equation  for  this  latter  reaction.     Is  it  a  reversible 
reaction  ? 

(b)  Repeat    (a)   using   separately  small   quantities   of    MnO2,   CuO, 
KC1O3,  and  PbO2. 

(c)  Place  a  piece  of  sodium  peroxide  (Na2O2),  about  the  size  of  a  pea, 
in  a  test  tube  half  filled  with  water.     Test  the  evolved  gas  as  in  a  (?). 
Equation  ? 

Hg  +  O  — >  HgO  +  30,600  cal. 

How  many  calories  of  heat  would  have  to  be  supplied  to  a  gram- 
molecular  weight  of  HgO  to  decompose  it  into  Hg  and  6  ? 

Experiment  II. — Laboratory  Source  of  Oxygen. 

Thoroughly  mix  10  grams  of  potassium  chlorate  (KC1O3)  with  exactly 
8  grams  of  MnO2  (manganese  di-  or  per-oxide)  ;by  means  of  a  folded  piece 
of  paper  introduce  the  mixture  into  the  large  test  tube  or  retort  (Fig.  8) 


which  you  will  find  in  the  drawer;  close  the  tube  with  a  rubber  cork  fitted 
with  delivery  tube;  clamp  the  tube  to  ring  stand  at  an  angle  of  about 
45°  so  that  delivery  tube  will  dip  beneath  the  surface  of  the  water  in  the 
pneumatic  trough ;  heat  lower  end  of  tube  cautiously,  and  fill  six  bottles 
with  the  gas  by  displacement  of  water;  place  bottles  in  upright  position 
and  cover  with  wet  paper.  Observe  color,  taste  and  odor  of  gas  as  it  is 


56  EXPERIMENTAL  CHEMISTRY. 

evolved;  remove  delivery  tube  from  pneumatic  trough;  when  tube  is  cold, 
add  water  to  contents  and  allow  to  soak  for  several  days  if  necessary 
to  remove  mixture ;  see  Exp.  VII.  Write  equation  for  reaction,  remember- 
ing that  MnO2  is  not  altered  by  the  reaction.  Proceed  with  following 
experiments. 

PROPERTIES    OF    OXYGEN.        "KINDLING      TEMPERATURE." 
OXIDATION — OXIDES. 

Experiment  III. — Oxidation  of  Carbon. 

(a)  Drop  a  small  piece  of  charcoal  into  one  of  the  bottles  containing 
oxygen;  immediately  recover  bottle.     Is  there  any  evidence  of  chemical 
action?     Heat  to  redness  another  small  piece  of  charcoal;  by  means 
of  a  pair  of  forceps,  deflagrating  spoon  or  a  piece  of  wire,  introduce 
glowing  charcoal  into  above  bottle  of  gas.     Results  ? 

(b)  Pour  a  little  clear  lime-water  (Ca(OH)2)  into  another  bottle  of 
the  gas;  place  your  hand  over  mouth  of  bottle  and  shake;  does  the  water 
change  in  color?     Now  introduce  a  piece  of  glowing  charcoal  as  per 
above  experiment;  when  charcoal  ceases  to  burn,  shake  bottle  as  before. 
Is  there  any  change  in  the  color  of  the  water?     Write  equations  repre- 
senting above  reactions.     If  97,000  calories  are  liberated  by  the  form- 
ation of  one  gram-molecule  of  CO2  from  free  O  and  free  C,  write  the 
"energy  equation"  for  the  reaction. 

The  "kindling  temperature"  of  a  substance  is  the  temperature  to 
which  it  must  be  raised  before  it  will  undergo  combustion.  It  is  definite 
for  a  particular  substance  but  varies  greatly  for  different  substances. 

Experiment  IV. — Oxidation  of  Phosphorus. 

Yellow  phosphorus  is  usually  moulded  into  small  cylinders.  Place 
one  of  these  pieces  under  water  and  then  by  means  of  a  knife  cut  off 
a  piece  about  the  size  of  a  small  pea.  Never  handle  P  with  the  hands, 
as  it  readily  catches  fire.  Cover  the  bottom  of  one  of  the  bottles  of  gas 
with  water;  test  water  with  litmus  paper  (blue);  place  phosphorus  in  a 
deflagrating  spoon  (Fig.  9)  and  ignite  it,  then  lower  it  into  bottle  of  gas. 
Results  ?  Test  water  in  bottle  with  blue  litmus  paper.  Write  equations 
for  reaction.  Was  energy  set  free  by  the  reaction  ? 

Experiment  V. — Oxidation  of  Sulphur. 

By  means  of  a  deflagrating  spoon,  lower  a  little  ignited  sulphur  into 
a  bottle  of  gas;  the  bottom  should  be  covered  with  water  and  tested  with 
litmus  paper  before  and  after  the  burning  of  the  sulphur.  Results? 
Was  energy  liberated?  Equations? 

Experiment  VI. — Oxidation  of  Iron. 

Lower  the  unwound  end  of  a  piece  of  "picture  cord"  (iron)  into  a 
bottle  of  the  gas  for  a  few  seconds.  Is  there  any  manifestation  of  chem- 


OXYGEN.  57 

ical  action?  Remove 'the  wire  and  dip  it  into  the  burning  sulphur  in 
the  deflagrating  spoon;  if  the  sulphur  on  the  wire  is  burning,  lower  wire 
into  bottle  (Fig.  10)  again.  Results  ?  The  bottom  of  the  bottle  should  be 
covered  previously  with  water  to  protect  it  from  the  molten  iron.  Write 
equation.  Why  was  it  necessary  to  burn  sulphur  on  the  end  of  the 
wire  ?  Was  energy  set  free  by  the  oxidation  of  the  iron  ? 

The  greatest  diversity  of  characteristics  exists  among  these  oxides, 
which  for  purposes  of  convenience  may  be  classed  either  as  metallic  or 
non-metallic  oxides.  Many  of  the  elements  are  capable  of  forming  a 
series  of  oxides;  nitrogen,  for  example,  forms  five  different  compounds 
with  oxygen.  When  more  than  one  oxide  is  formed,  the  suffix  ous  or  ic 
is  attached  to  the  name  of  the  metal  accordingly  as  the  oxide,  relative  to 
the  metal,  contains  a  lesser  or  greater  amount  of  oxygen,  i.e.,  these 
endings  indicate  the  relative  degrees  of  oxidation  which  the  metals  have 


FIG.  9.  FIG.  10. 

undergone,  thus  iron  (ferrum)  forms  two  oxides,  ferrous  oxide  (FeO), 
and  ferric  oxide  (Fe2O3) ;  mercury  possesses  the  same.property  of  forming 
two  oxides,  mercurous  (Hg2O)  and  mercuric  (HgO)  oxides;  copper 
forms  cuprous  oxide  (Cu2O),  and  cupric  oxide  (CuO). 

Inasmuch  as  oxygen  is  ordinarily  bivalent,  it  is  obvious  that  a  metal 
frequently  shows  a  varying  valency  as  represented  by  its  ous  and  ic 
compounds. 

In  the  preceding  experiments  involving  "oxidation,"  it  was  readily 
noted  that  the  reactions  were  exothermic,  i.e.,  the  "heat  of  formation" 
of  the  respective  oxides  was  positive;  it  follows  that  the  energy  (usually 
heat)  necessary  to  decompose  ("heat  of  decomposition")  an  oxide  is 
equal  to  the  energy  (usually  light  and  heat)  liberated  during  its  formation. 

This  may  be  made  clear  by  the  use  of  an  analogy.  As  mentioned 
previously,  substances  capable  of  directly  combining,  possess  a  definite 
amount  of  chemical  energy  which  may  be  likened  unto  the  potential 
energy  which  a  stone  possesses  when  it  is  raised  above  the  ground;  when 
the  substances  combine,  all  or  portions  of  the  chemical  energy  of  the 
factors  is  transformed  into  say  heat  and  light  energy,  just  as  the  potential 
energy  of  the  stone  is  converted  into  kinetic  energy  by  permitting  the 
stone  to  fall ;  now  it  will  require  just  as  much  energy  to  raise  the  stone  to 


58  EXPERIMENTAL  CHEMISTRY. 

its  original  elevated  position  as  was  freed  during  its  descent;  just  so 
with  chemical  compounds,  the  "heat  of  decomposition"  is  necessarily 
equal  to  the  "heat  of  formation."  If  combination  is  attended  with  the 
evolution  of  much  energy  (heat  and  light)  the  oxides  will  be  stable, 
and  vice  versa. 

Although  the  "heat  of  combustion"  of  carbon  in  oxygen  with  forma- 
tion of  CO,  can  not  be  directly  measured,  it  can  be  calculated  from 
calorimetric  measurements  indicated  by  the  following  equations: 

C      +  2O  —  CO2  +  97,000  cal. 

and         CO  +     O  —  CO2  +  68,000  cal. 

subtracting,    C      +     O— >  CO    +  29,000  cal. 

Experiment    VII.— Catalytic  Action  of  Manganese  Dioxide  in  Exp.  II. 

This  experiment  is  a  continuation  of  Exp.  II.  Half  fill  the  test  tube 
with  water;  heat  gently  in  an  endeavor  to  dissolve  contents  of  tube;  be 
careful  to  avoid  losing  any  of  the  solid  material;  pour  the  clear  super- 
natant liquid  upon  a  filter;  avoid  pouring  much  of  the  solid  matter  upon 
the  filter;  half  fill  the  tube  again  with  water  and  heat  to  boiling;  decant 
the  fluid  as  before  upon  the  filter;  repeat  above  operation  four  or  five 
times,  then  pour  contents  of  tube  on  filter;  rinse  out  tube  thoroughly, 
and  wash  material  upon  filter  with  hot  water;  spread  filter  and  contents 
upon  wire  gauze  and  heat  gently  until  paper  is  dry;  place  paper  upon 
glass  plate  and  scrape  black  powder  into  a  weighed  crucible ;  fold  filter 
paper  and  place  it  in  crucible ;  heat  crucible  with  low  flame  until  contents 
are  perfectly  dry,  and  filter  paper  has  ignited  and  burned  to  ash ;  avoid 
heating  crucible  to  redness;  cool  and  weigh;  ignore  weight  of  ash  of 
filter  paper;  what  is  the  weight  of  the  black  powder?  Does  this 
powder  resemble  MnO2  ?  Do  your  results  justify  the  conclusion  that 
WnO-j  was  not  altered  in  the  reaction  ?  Add  a  few  drops  of  the  filtrate  to 
a  cubic  centimeter  of  silver  nitrate.  Results  ?  Add  a  few  drops  of  a 
potassium  chlorate  solution  to  a  cubic  centimeter  of  silver  nitrate.  Re- 
sults ?  Did  the  KC1O3  undergo  a  change  in  Exp.  II  ? 

2KC1O3  +  2MnO2— »  2KMnO4  +  2C1  +  2O 
2KMnO4  —  K2MnO4  +  MnO2  +  2O_ 
K2Mn04  +  2C1—  2KC1  +  MnO2  +  2!) 
adding,  2KC1O3  +  2MnO2—  2KC1  +  2MnO2  +  6O(3O2). 

A  substance  which  alters  the  speed  of  a  chemical  change  apparently 
by  its  mere  presence  and  contact,  without  undergoing  any  permanent 
change,  is  called  a  catalytic  agent  or  simply,  a  catalyser.  It  may  be  either 
a  gas,  liquid  or  solid.  Pieces  of  platinum  foil  are  frequently  used  to 
catalyse  a  reaction.  The  process  itself  is  called  catalysis. 

"To  obtain  a  picture  of  the  way  in  which  a  catalyser  acts,  imagine  a 
wheel-work  in  which  the  axles  move  with  great  friction,  as  a  result  say  of 
the  oil  having  become  thick,  and  which  therefore  runs  down  only  very 


OXYGEN.  59 

slowly.  If  a  little  fresh  oil  be  placed  on  the  axles  the  wheel-work  forth- 
with runs  down  much  more  quickly,  although  the  available  tension  of 
the  spring  (which  corresponds  to  the  work  available  from  chemical 
reaction)  is  in  no  way  altered  by  the  oil.  The  action  of  a  catalyser  may 
be  compared  with  that  of  the  oil  in  this  respect,  and  also  with  respect 
to  the  fact  that  the  oil  is  not  used  up  in  acting." — Ostwald. 

It  should  be  understood  that  a  catalyser  does  not  initiate  a  reaction; 
it  is  not  the  cause  of  a  reaction;  it  merely  increases  or  diminishes  the 
speed  of  a  given  chemical  reaction. 

Experiment  VIII. — Preparation  and  Properties  of  Ozone  (O3). 

Place  two  or  three  pieces  of  yellow  phosphorus,  about  the  size  of  a 
"  playing"  marble,  in  a  bottle;  add  a  sufficient  quantity  of  water  to  cover 
bottom  of  bottle,  but  not  any  more  than  will  be  necessary  to  half  cover 
the  pieces  of  phosphorus;  four  or  five  drops  of  a  solution  of  potassium 
dichromate  should  be  added  to  the  water;  cork  the  bottle.  Now  prepare 
a  piece  of  test-paper  as  follows:  dip  a  strip  of  filter  paper  into  a  starch 
emulsion  to  which  has  been  previously  added  a  few  drops  of  a  solution 
of  potassium  iodide  (KI);  let  the  strip  of  paper  "drip"  for  a  few  minutes, 
then  suspend  it  from  the  cork  in  the  bottle  so  that  paper  is  within  a  centi- 
meter of  the  phosphorus;  allow  the  apparatus  to  stand  for  one  laboratory 
period,  then  examine  paper.  Note  the  odor  of  the  gas  in  the  bottle. 
Results?  Equations?  The  phosphorus  should  be  returned  to  the 
supply  bottle. 

Note. — A  good  test  paper  for  above  experiment  is  made  by  dipping 
a  piece  of  red  litmus  paper  into  a  solution  of  KI.  When  the  O3  acts  upon 
the  KI,  the  iodine  is  set  free  and  KOH  acts  upon  the  red  litmus  paper 
turning  it  blue.  Satisfactory  results  are  frequently  secured  by  placing 
phosphorus  in  a  dry  flask  which  is  then  closed  with  a  cork. 

A  quantitative  study  of  the  conversion  of  oxygen  (O2)  into  ozone 
(O3)  reveals  that  3  volumes  of  the  former  are  required  to  form  2  volumes 
of  the  latter,  and  that  the  reaction  is  endothermic,  i.e.,  absorbs  heat 
energy. 

(1)  02  +  O  —  03  —  32,900  cal. 

(2)  O3  —  O2  +  O   +  32,900  cal. 

At  a  temperature  of  about  275°  C.  the  above  reaction  (i)  is  reversed. 
Ozone  is  a  gas  of  blue  color,  odor  like  dilute  chlorine,  density  24  (H2  =  i); 
molecular  weight,  47.9;  boils  at  — 119°  C.  Ozone  is  much  more  soluble 
than  oxygen  in  water — 100  volumes  of  water  at  12°  C.  dissolves  50 
volumes  of  ozone  under  a  pressure  of  one  atmosphere.  The  chemical 
properties  of  oxygen  and  ozone  are  similar,  save  that  the  latter  is  much 
more  active. 

It  is  evident  that  oxygen  and  ozone  are  identical  so  far  as  the  kinds 


60  EXPERIMENTAL  CHEMISTRY. 

of  matter  of  which  they  are  composed  is  concerned.  It  seems  to  be 
equally  as  evident  that  their  different  properties  are  due  to  the  fact  that 
they  differ  in  their  respective  energy  contents.  When  ozone  is  used 
as  an  oxidizer,  32,900  more  calories  are  liberated  than  when  an  equal 
weight  of  oxygen  is  used,  hence  the  activity  of  ozone  as  an  oxidizer. 
It  is  to  be  noted  that  the  heat  developed  by  the  passing  of  ozone  into 
ordinary  oxygen  is  not  contained  in  the  ozone,  as  heat  energy  but  in  the 
form  of  chemical  energy. 

Ozone  and  oxygen  are  allotropic  modifications  of  the  same  element. 

Experiment  IX. — (Quant. )  Determination  of,  (a)  the  Density  of  Oxygen, 
(b)  the  Volume  of  Oxygen  Liberated  from  a  Given  Weight  of  Potassium 
Chlorate,  (c)  the  Weight  of  the  Potassium  Chloride  Formed. 

(a)  Note. — i  gram  of  KC1O3  liberates  about  290  cm. 3  of  O2. 

Place  a  gas  burette,  wrhich  is  connected  with  a  leveling  bulb,  in  a  water 
jacket;  suspend  a  thermometer  in  the  water  jacket;  raise  leveling  bulb 
until  water  in  burette  stands  at  top  of  capillary  tube  on  upper  end  of 
burette.  Clean  and  dry  and  weigh  a  hard- glass  test  tube;  record  weight; 
introduce  into  the  tube  .12  grm.  of  KC1O3  which  has  been  previously 
pulverized,  and  dried  in  an  air  bath  (100°  C.)  for  one  hour.  Connect  test 
tube  with  burette  by  means  of  rubber  tubing  and  a  piece  of  capillary 
glass  tubing  bent  into  such  form  as  will  permit  test  tube  to  be  clamped  at 
an  angle  of  about  45  °.  All  joints  must  be  air  tight.  When  making  the 
connections  be  sure  that  no  water  is  forced  into  the  glass  tubing.  Adjust 
the  leveling  bulb  so  that  \vater  in  it  and  in  burette  are  at  the  same  level. 
Read  burette  and  thermometer,  and  record  readings.  Holding  burner 
in  hand  gently  warm  the  test  tube;  oxygen  will  be  liberated  and  pass 
over  into  the  burette.  Always  keep  the  water  in  the  leveling  bulb  about 
3  or  4  cm.  below  the  water  in  the  burette.  Continue  to  heat  tube  until 
gas  ceases  to  be  evolved,  then  remove  burner  and  allow  the  apparatus 
to  assume  its  original  temperature,  then  bring  the  water  in  burette  and 
bulb  to  same  level.  Read  burette,  thermometer  and  barometer,  and 
record  readings.  From  tables  in  Appendix  get  vapor  tension  (a)  of 
water  at  the  temperature  of  the  water  jacket.  Reduce  the  volume 
of  oxygen  to  standard  conditions. 

Vol.  x  273  x  (Bar.  P.  — a) 


(273  +  t)  x  760 

Remove  the  test  tube;  wipe  and  weigh  it;  what  is  the  weight  of  the 
KC1?  The  difference  between  the  weight  and  the  original  weight  of 
the  KC1O3  is  the  weight  of  the  oxygen  evolved.  Calculate  the  weight 
of  a  cm.3  of  oxygen  at  standard  conditions.  The  weight  of  a  liter? 
Calculate  the  weight  and  volume  of  oxygen  which  should  have  been 
liberated  from  the  KC1O3  on  a  purely  theoretical  basis. 


OXYGEN. 


6T 


Experiment  X. — (Quant.)  Optional.  Determination  of  the  Percentage  of 
Oxygen  in  the  Air. 

Note. — The  general  theory  of  this  experiment  is  to  isolate  a  given 
volume  of  air  under  known  conditions,  then  bring  this  volume  of  air  into 
intimate  contact  with  a  substance  which  will  absorb  the  oxygen,  after 
which  the  volume  of  air  is  measured  under  the  original  conditions, 
and  its  loss  in  volume  noted. 

Use  the  measuring  apparatus  assembled  for  Exp.  IX  or  Hempel 
burette  (Fig.  n).  Attach  a  piece  of  capillary  rubber  tubing  to  the 
upper  end  of  the  gas  burette;  the  tubing  should  be  about  7  cm.  long 
and  firmly  wired  to  the  burette;  place  a  Mohr  pinchcock  upon  the 
rubber  tubing  close  to  the  end  of  the  burette;  open  pinchcock;  raise 
leveling  bulb  until  water  in  burette  stands  at  top  of  capillary  tubing, 
then  lower  bulb  until  about  40  cm.3  of  air  have  been  drawn  into 
burette;  close  pinchcock;  allow  gas  to  stand  for  five  minutes;  bring 
water  in  burette  and  bulb  to  the  same  level;  read  thermometer  sus- 


FIG.  ii. — Hempel  Burette. 
(Smith  and  Keller.) 


FIG.  12. — Hempel  Compound  Pipette. 
(Smith  and  Keller.) 


pended  in  water  jacket,  read  barometer,  read  burette,  and  make  a  record 
of  readings;  now  a  compound  gas  pipette  (any  desirable  form,  Fig.  12) 
containing  an  alkaline  solution  of  pyrogallol,*  should  be  connected 
with  the  gas  burette  by  means  of  a  short  piece  of  capillary  glass  tubing, 
the  pipette  should  be  closed  by  rubber  tubing  and  a  pinchcock;  when  the 
pipette  and  burette  are  properly  connected  slightly  raise  the  leveling 
bulb,  then  open  both  pinchcocks  and  by  gradually  raising  the  bulb  the 
air  will  be  forced  over  into  the  pipette;  raise  bulb  until  water  in  burette 
stands  at  top  of  rubber  tubing  on  burette;  close  both  pinchcocks;  dis- 
connect pipette  and  burette;  shake  pipette  for  20-25  min-5  connect 
pipette  with  burette;  gradually  lower  bulb,  at  the  same  time  open  the 
pinchcocks;  continue  to  lower  bulb  until  all  the  gas  is  drawn  from  pipette 

*The  alkaline  "  pyro  "  solution  is  prepared  as  suggested  by  Professor  Baxter  of 
Harvard,  as  follows — 

One  gram  of  pyrogallol  to  4  cm. 3  of  water;  i  gram  of  KOH  to  i  cm. 3  of  water.  The 
solution  prepared  in  above  ratio  must  be  mixed  out  of  contact  with  the  air;  the  pyro 
solution  is  first  introduced  into  the  pipette  after  which  the  KOH  solution  follows; 
shake  until  solutions  are  thoroughly  mixed. 


62  EXPERIMENTAL  CHEMISTRY. 

and  the  "pyro"  solution  is  drawn  up  to  the  top  of  rubber  tubing  on 
pipette;  close  pinchcocks,  be  sure  that  both  are  closed  tightly;  allow 
gas  in  burette  to  stand  for  5  min.;  bring  water  in  burette  and  bulb  to 
same  level;  read  burette,  thermometer,  barometer  and  record  readings. 
Repeat  the  operation  until  no  further  decrease  in  volume  occurs.  Reduce 
both  the  original  and  final  volumes  of  air  to  standard  conditions.  What 
is  the  total  change  in  volume?  'Calculate  the  percentage  of  oxygen  in 
the  air  from  your  measurements. 

A  principle  of  great  importance,  especially  in  connection  with  the 
measurement  of  the  volume  of  gases,  is  incorporated  in  Dalton's  Law 
of  Partial  Pressures: 

The  pressure  exerted  by  a  mixture  oj  a  gas  and  a  vapor,  of  two  vapors, 
or  oj  two  gases,  is  equal  to  the  sum  oj  the  pressures  which  each  would  exert 
if  it  alone  occupied  the  whole  space  afforded  to  the  mixture. 

This  law  is  not  absolutely  exact.     It  may  be  stated  in  another  form: 

The  volumes  oj  two  or  more  gases  in  a  volume  oj  the  mixture  are  pro- 
portional to  their  respective  pressures. 

Another  of  Dalton's  Laws  is: 

The  pressure  exerted  by,  and  the  quantity  oj  a  vapor  which  saturates  a 
given  space  are  the  same  for  the  same  temperature  whether  this  space  is 
filled  by  a  gas  or  is  a  vacuum. 

PROBLEMS. 

i. — How  many  grams  of  sulphur  will  exactly  combine  with  32  grm. 
of  oxygen  to  produce  SO2? 

2. — What  weight  of  oxygen  could  be  obtained  from  10  grm.  of  KC1O3 
if  the  latter  contains  10  per  cent,  of  an  impurity?  Now  many  liters,  if 
a  liter  of  oxygen  weighs  1.428  grm.  ? 

3. — What  is  the  "heat  of  formation"  of  SO2  if  2  grm.  of  sulphur  in 
burning  to  sulphur  dioxide  develops  4440  calories?  Write  the  "energy" 
equation. 

4. — The  specific  heat  of  copper  is  .092.  How  many  calories  of  heat 
are  liberated  when  100  grm.  of  copper  cool  from  75°  C.  to  50°  C.  ? 

5. — Is  the  heat  liberated  during  formation  of  an  oxide,  necessarily 
an  accurate  measure  of  the  chemical  energies  (affinity)  of  the  reacting 
substances  ?  Explain. 

6. — Reduce  200  cm.s  of  gas  at  20°  C.  and  730  mm.  to  o°  and  760  mm. 

7. — Reduce  45  cm.3  of  gas  at  —  20°  C.  and  770  mm.  to  o°and  760  mm. 

8. — Reduce  40  cm.3  of  gas  collected  over  water  at  20°  C.  and  a  baro- 
metric reading  of  730  mm.  to  o°  C.  and  760  mm. 

9. — A  gas  globe  when  full  of  air  weighed  55.06  grm.;  full  of  water  at 


OXYGEN.  63 

20°  C.  it  weighed  309.66  gram.  The  globe  was  "  exhausted  "  (air  removed) 
and  carefully  weighed  by  means  of  a  counterpoise;  it  was  then  filled  with 
a  gas  at  19.8°  C.  and  761.4  mm.,  when  it  weighed  .469  grm.  more  than 
when  exhausted.  What  is  the  weight  of  a  liter  of  this  gas  at  standard 
conditions  ? 

10. — What  is  the  molecular  weight  of  the  gas  used  in  above  experi- 
ment? 


CHAPTER  IX. 
HYDROGEN. 

Symbol— H.     At.  Wt.,  1.0075. 

Hydrogen  is  a  light  colorless  gas,  first  described  as  a  form  of  air 
(phlogiston)  by  Cavendish,  in  1776,  although  it  had  undoubtedly  been 
observed  previously  by  Paracelsus  and  Boyle.  Lavoisier  gave  the  ele- 
ment the  name,  hydrogene  (to  produce  water).  It  is  univalent. 

Experiment  I. — Methods  for  Preparing  Hydrogen. 

(a)  Electrolysis  of  water  (H2O).     Recall,  or  repeat  this  experiment 
as  previously  performed.     At  which  pole  was  the  hydrogen  liberated? 
What  test  for  hydrogen  was  used  ? 

(b)  I. — Decomposition  of  water  by  means  of  metals.     Wrap  a  piece 
of  sodium  or  potassium  about  the  size  of  a  pea  in  a  small  piece  of  filter 
paper;  place  it  in  a  small  wire  cage  and  dip  it  under  the  surface  of  the  water 
in  the  pneumatic  trough  (Fig.  13);  hydrogen  gas  will  be  evolved,  which 
collect  in  a  test  tube  by  displacement  of  water;  when  action  has  ceased 
place  your  thumb  over  the  mouth  of   tube  and  hold  tube  in  upright 
position;  place  a  glowing  splinter  in  the  mouth  of  tube  for  a  second. 
Results?     Bring  a  lighted  match  to  the  mouth  of  the  tube.     Results? 
Does  the  hydrogen  support  combustion  so  far  as  the  glowing  match  is 
concerned  ?     When  raised  to  a  given  temperature  does  it  combine  with  a 
component  of  the  air? 

H20    +  K->KOH  +  H 
Water  Potassium 

Hydroxide, 
Potassium,  Hydrogen. 

H2O  +  Na  —  NaOH  +  H. 

(b)  II. — In  the  above  experiment  all  of  the  hydrogen  of  the  molecule 
of  water  was  not  displaced  by  the  metal;  this  may  be  accomplished  by 
another  process.  Pulverize  and  mix  thoroughly  and  quickly  about  i 
grm.  of  sodium  hydroxide  (NaOH)  and  2  grm.  of  zinc  dust;  introduce 
mixture  into  a  hard  glass  test  tube;  heat  tube  and  test  for  hydrogen  gas. 
Results  ? 

Zn  +  2NaOH  —  Zn(ONa)2  +  2H. 

(b)  III. — Decomposition  of  water  by  magnesium.  Shake  a  gram  of 
magnesium  powder  into  the  bottom  of  a  test  tube  and  add  10  cm. 3  of  water; 

64 


HYDROGEN.  65 

add  a  few  cm. 3  of  magnesium  chloride  (this  does  not  take  any  part  in  the 
reaction,  merely  dissolving  the  magnesium  oxide  which  is  formed  on  the 
surface  of  the  metal).  Heat  the  tube  and  its  contents;  test  for  hydrogen 
gas.  Results  ?  Equation  ? 

(c)  Displacement  of  Hydrogen  from  Acids  by  Metals. 

(c)  I. — To  a  small  piece  of  each  of  the  following  metals,  placed  in  sep- 
arate test  tubes  add  5-10  cm.3  of  hydrochloric  acid,  HC1  (shelf-reagent); 
granulated  zinc,  magnesium  (ribbon),  iron  (filings).  Test  each  tube 
for  the  presence  of  hydrogen.  Notice  whether  the  test  tubes  become  warm 
as  the  reaction  continues.  What  is  the  effect  of  heating  the  tubes  ?  Re- 
cord all  data.  Assume  that  all  of  the  above  metals  are  bivalent;  repre- 
sent the  reactions  by  "energy  equations"  using  the  word  "heat"  instead 
of  calories.  Does  the  final  "system"  contain  more  or  less  energy  than 
the  original  one?  Why? 


FIG.  13. 

(c)  II. — Repeat  (c)  I,  using  cold  dilute  sulphuric  acid  (H2SO4)  instead 
of  HC1.  This  dilute  acid  may  be  prepared  by  pouring  i  volume  of  con- 
centrated H2SO4  (shelf-reagent)  into  4  volumes  of  water;  allow  mixture 
to  cool.  Add  a  few  drops  of  copper  sulphate  solution  to  test  tubes  con- 
taining iron  and  zinc  respectively.  What  is  the  effect  ?  Similar  results 
are  secured  by  placing  a  piece  of  platinum  foil  or  wire  in  contact  with 
the  zinc.  This  may  be  verified  by  the  student  by  using  a  separate  test 
tube  containing  zinc  and  dilute  H2SO4.  After  the  zinc  has  dissolved 
completely,  filter  the  solution  which  remains;  evaporate  the  filtrate  until 
a  thin  film  appears  on  the  surface  when  the  solution  cools.  Are  crystals 
deposited  ?  Does  it  resemble  zinc  or  H2SO4  ?  Repeat  this  process  using 
the  iron  solution.  Write  equations  as  in  (c)  I.  Can  you  explain  the 
action  of  the  platinum  or  the  copper  sulphate  ?  Did  the  platinum  lose  its 
identity  ? 


66  EXPERIMENTAL  CHEMISTRY. 

(c)   III. — Effect  of  Surface  on  Speed  of  Reaction. 

Try  the  action  of  dilute  H2SO4  on  zinc  dust.  A  small  quantity  of  the 
latter  placed  in  the  bottom  of  t.t.  will  be  sufficient.  Compare  results 
with  action  of  the  acid  on  "granulated"  zinc.  Explain. 

(c)  IV. — Compare  the  interaction  of  acetic  acid  (HC2H3O2)  and  zinc 
with  foregoing  acids  and  zinc.  Repeat  above  using  concentrated  H2SO4. 

The  interaction  of  certain  metals  with  water  is  identical  in  principle 
with  the  interaction  of  certain  acids  and  metals;  the  metallic  substance 
hydrogen,  is  attached  in  both  cases  to  a  negative  element  or  radical, 
thus,  H.OH,  H.C1,  H2.SO4;  the  hydrogen  (metal)  is  then  displaced  by 


FIG.  14. 

certain  other  metals,  (not  all  metals,  for  ex.  Au.,  Ag.,  Pt.)  which  possess 
a  greater  chemical  energy  when  in  contact  with  the  above  mentioned 
negative  groups.  A  reaction  then  occurs  whereby  a  new  system  is  es- 
tablished which  contains  less  energy  than  the  original  one. 

Zn  +  H2SO4  —  ZnSO4  +  2H. 

The  " system"  represented  by  the  bodies  to  the  lejt  of  the  arrow  in  the 
above  chemical  equation  obviously  contains  more  chemical  energy  than 
the  system  represented  by  those  to  the  right,  as  evidenced  by  the  fact  that 
heat  was  given  out  during  the  reaction. 

Experiment  II. — Laboratory  Source  of  Hydrogen. 

Assemble  a  gas  bottle  (generating  flask);  use  the  heavy  glass  flask 
(Fig.  14)  of  about  500  cm.  3  capacity;  fit  flask  with  a  rubber  cork  perforated 
with  two  holes;  through  one  of  the  holes  pass  a  thistle  tube;  in  the  other 
fit  a  delivery  tube  with  its  arms  bent  at  90°;  connect  delivery  tube  of 
generator  with  a  washing  bottle  ;half-filled  with  concentrated  H2SO4; 
washing  bottle  should  be  provided  with  delivery  tube  which  dips  beneath 
surface  of  water  in  pneumatic  trough.  Place  enough  granulated  zinc 


HYDROGEN. 


67 


in  the  flask  to  cover  the  bottom  completely;  add  dilute  H2SO4  through 
thistle  tube  (i  of  acid  to  4  of  H2O)  as  needed  to  secure  a  rapid  evolution 
of  gas;  be  sure  that  end  of  thistle  tube  dips  beneath  the  liquid  in  flask. 
If  gas  is  not  evolved  rapidly  add  a  few  drops  of  copper  sulphate  solution 
through  thistle  tube.  Collect  the  gas  in  a  test  tube  by  displacement  of  water 
until  the  hydrogen  burns  quietly  in  the  test  tube  when  a  flame  is  applied 

to  it,  then  collect  several  test  tubes  of  the 
gas  and  proceed  with  a  study  of  the  proper- 
ties of  hydrogen. 

Note. — The  test  tubes  containing  the 
gas  must  remain  inverted  with  mouths 
under  water.  Proceed  with  Experiments 
IV  and  V. 

Experiment  III. — Properties  of 
Hydrogen. 

(a)  Remove  one  of  the  test  tubes  con- 
taining   hydrogen   from    the   pneumatic 
trough  and  observe  whether  the  gas  has 
any  color,  odor,  etc. 

(b)  Allow  one  of  the  test  tubes  filled 
with  hydrogen  to  stand  mouth  upward 
and  uncovered  for  a  minute.      Apply  a 
flame.     Results?     Explain.     Is   the  gas 
apparently  lighter  or  heavier  than  air? 

(c)  Place  your  thumb  over  the  mouth 
of  a  tube  of  gas;  hold  it  vertically,  mouth 
upward;  remove  the  thumb  and  immedi- 
ately bring  the  mouth  of  a  test  tube  con- 
taining air,  down  over  the  mouth  of  the 
tube  of  hydrogen,  in  such  a  manner  as  to 
prevent  the  gas  escaping  from  the  two 
tubes,  but  will  permit  them  to  "diffuse" 
into  one  another;  hold  them  in  this  posi- 
tion for   i   min. ;    apply  a  flame  to  the 

mouth  of  each  tube.     Results  ?     Explain. 

Tabulate  the  properties  of  hydrogen  as  revealed  by  above  experiments. 

Experiment  IV. — Effusion  of  Hydrogen  Through  a  Porous  Medium. 

As  relatively  large  quantities  of  hydrogen  are  required  for  this  experi- 
ment (Fig.  15),  it  can  be  performed  to  advantage  by  the  instructor  or 
assistant. 

Law  of  Effusion,  of  Graham  and  Bunsen. — The  velocities  oj  effusions 
of  gases  are  inversely  proportional  to  the  square  roots  oj  the  densities. 
Explain  on  basis  of  "  Kinetic  Theory  of  Gases." 


FIG. 


68  EXPERIMENTAL  CHEMISTRY. 

Experiment  V. — Synthesis  of  Water  by  Burning  Hydrogen  in  Air. 

Draw  a  piece  of  hard  glass  tubing  out  into  the  form  of  a  fine  jet.  At- 
tach this  jet  to  the  delivery  tube  of  the  washing  bottle  (drying  agent  in 
this  case);  permit  jet  to  dip  beneath  surface  of  water  in  pneumatic  trough; 
fill  a  test  tube  with  the  gas  by  the  usual  method;  raise  jet  out  of  the  water; 
apply  flame  to  test  tube  of  gas;  quickly  pass  the  tube  over  the  jet  and 
immediately  withdraw  it;  if  gas  is  pure  it  burns  comparatively  slowly  so 
that  when  tube  is  passed  over  jet,  the  gas  issuing  from  latter  is  ignited; 
if  gas  is  impure,  i.e.,  diluted  with  oxygen,  it  explodes  when  flame  is  ap- 
plied to  test  tube.  Always  light  the  gas  jet  by  above  method.  Place  over 
the  burning  jet  a  cold  dry  bottle  or  let  flame  impinge  upon  a  cold  dry 
glass  plate  (Fig  16).  Does  water  form  upon  the  sides  of  the  glass?  Is 
the  combination  of  hydrogen  and  oxygen  to  form  water  an  endo-  or 


FIG.  1 6.— (Smith  and  Keller). 

exothermic  reaction?  Your  reasons?  Explain.  Equation?  When  a 
lamp  in  a  cold  room  is  first  lighted  moisture  appears  on  inside  of  chimney. 
Explain.  When  bottles  of  varying  sizes  are  placed  over  a  burning  jet, 
a  musical  note  is  frequently  emitted;  it  has  been  called  the  "singing 
flame."  Explain. 

H2  +  O  —  H2O,  Aq  +  68,400  cals. 
Experiment  VI. — Hydrogen  and  Air  Form  an  Explosive  Mixture. 

The  Hydrogen  Cannon.  Into  a  small  tin  vessel  closed  at  one  end 
with  a  cork,  and  containing  a  small  opening  at  the  other  end,  introduce 
a  stream  of  hydrogen  for  a  few  seconds;  apply  a  flame  to  the  small  open- 
ing after  " pointing"  cork  toward  the  ceiling.  Explain.  A  mixture  of 
hydrogen  and  oxygen  gases  is  called  "detonating  gas."  Soap-bubbles 
filled  with  this  gas  will  explode  when  a  flame  is  applied,  with  a  report  like 
a  gun. 

Experiment  VII. — Occlusion  of  Hydrogen.  Catalytic  Action  of 
"Spongy  Platinum." 

Prepare  a  piece  of  " platinized  asbestos"  by  dipping  a  small  piece  of 
" sheet  asbestos"  in  a  solution  of  platinic  chloride,  and  then  in  an  ammo- 
nium chloride  solution;  heat  asbestos  in  the  hottest  portion  of  the  flame 
of  a  Bunsen  burner  or  the  blast-lamp;  finely  divided  platinum  is  depos- 
ited in  the  asbestos  by  this  process.  When  asbestos  has  cooled  slightly 


HYDROGEN.  69 

hold  it  by  means  of  forceps  or  wire  so  that  a  current  of  hydrogen  may 
pass  over  it.  Does  the  asbestos  begin  to  glow?  Is  the  gas  eventually 
raised  to  the  "  kindling  temperature  "  ?  Repeat,  using  natural  gas  or  il- 
luminating gas. 

"Apparently  connected  with  the  catalytic  action  of  platinum  is  its 
property  of  dissolving  large  quantities  of  different  gases,  especially  hy- 
drogen." "The  hydrogen  thereby  increases  enormously  in  reactivity 

"     "  It  must  not  be  supposed  that  the  chemical  affinity  or  the 

chemical  potential  of  the  hydrogen  is  changed;  such  an  assumption,  which 
is  certainly  very  often  made,  would  be  a  contradiction  of  the  fundamental 
laws  of  the  theory  of  energy."  "The  cause  of  the  changed  action  of  the 
platinum  lies  rather  in  the  'acceleration'  of  the  reactions  of  hydrogen  and  is, 

o 


FIG.  17. — (Smith  and  Keller.) 

therefore,  a  catalytic  action.  Gaseous  hydrogen  reacts  so  slowly  at  the  or- 
dinary temperature  that  it  appears  like  an  indifferent  substance,  and 
from  the  fact  that  in  the  presence  of  platinum  the  reaction  becomes  visible 
in  a  short  time,  while  otherwise  it  would  require  hours  or  perhaps  years, 
the  view  has  arisen  that  there  is  a  change  of  the  chemical  potential. "- 
Ostwald.  The  finely  divided  platinum,  exposing  much  surface,  is 
simply  a  catalytic  agent. 

Experiment  VIII. — Reduction  of  Cupric  Oxide — Oxidation  of  Hydro- 
gen— Synthesis  of  Water. 

Assemble  hydrogen  generator;  connect  it  with  washing  bottle  (Fig.  17) 
which  should  be  half  filled  with  H2SO4  (cone.),  that  the  hydrogen  gas  shall 
be  dry  after  bubbling  through  it;  you  will  find  in  the  drawer  a  piece  of  hard 
glass  tubing  i  cm.  in  diam.  and  18  cm.  long;  fit  into  one  end  a  rubber 
stopper  perforated  by  a  single  hole;  by  means  of  glass  tubing  connect  this 
end  of  the  tube  with  the  washing  bottle — the  hard  glass  tube  should  be  in  a 
nearly  horizontal  position;  the  entire  system  may  now  be  flooded  with 
hydrogen;  while  wraiting  for  all  of  the  air  to  be  removed  from  system, 
spread  i  gram,  of  cupric  oxide  (CuO)  over  the  bottom  of  a  small  porcelain 
boat  which  should  then  be  placed  in  the  middle  of  the  horizontal  tube; 
the  portion  of  the  tube  beneath  the  boat  should  be  warmed  gently  at  first, 
and  later,  quite  strongly.  Does  moisture  condense  on  cold  portions  of  the 
tube  ?  Is  there  any  evidence  that  the  CuO  has  undergone  any  change  ? 


70  EXPERIMENTAL  CHEMISTRY. 

Remove  heat;  draw  boat  out  of  tube  (if  water  comes  in  contact  with  hot 
tube  it  will  crack) ;  notice  the  color  of  the  powder.  Has  metallic  copper 
been  deposited  ?  Write  equations  representing  above  chemical  reactions. 

Note. — In  above  experiment  the  boat  may  be  dispensed  with;  place  the 
CuO  in  the  tube.  If  time  permits  the  reddish-brown  powder  may  be 
reoxidized  by  passing  a  slow  current  of  oxygen  over  it  when  heated  to 
redness.  Equation  ? 

Experiment  IX. — The  Nascent  State.     Nascent  Hydrogen. 

To  10  cm.3  of  a  dilute  solution  of  potassium  permanganate,  KMnO4, 
add  an  equal  volume  of  dilute  sulphuric  acid — after  shaking  divide  the 
solution  into  two  parts.  To  one  portion  in  a  test  tube  add  a  little  zinc  dust ; 
test  the  evolved  gas  with  a  flame.  What  is  the  effect  of  freshly  liberated 
hydrogen,  i.e.,  at  instant  of  liberation  or  birth,  upon  the  colored  solution? 

Pass  a  stream  of  hydrogen  gas  through  the  other  portion  of  the  solution. 
Results  ?  Explain. 

Experiment  X. — Properties  of  Hydrogen  Dioxide  (H2O2). 

(a)  From   the   "side-shelf"   reagents   procure   5-10   cm.3    of    H2O2. 
Observe  its  properties,  odor,  color,  reactive  properties  toward  red  and 
blue  litmus  paper,  etc. 

(b)  Pour  a  few  drops  of  H2O2  upon  a  piece  of  "test-paper,"  or  better, 
upon  a  little  starch  paste  containing  a  few  drops  of  a  KI  solution.     Re- 
sults ?     Dip  a  strip  of  red  litmus  paper  into  a  dilute  solution  of  potassium 
iodide  (KI);  then  add  two  or  three  drops  of  H2O2  to  litmus  paper;  explain 
why  litmus  paper  turns  blue.     Write  equations  showing  nature  of  above 
reactions. 

(c)  Heat  5-6  cm.3  of  H2O2  to  boiling;  test  for  evolved  oxygen;  add  a 
pinch  of  powdered  MnO2  or  any  powdered  metal,  preferably  the  former, 
to  the  H2O2  in  the  test  tube;  test  for  oxygen.     The  MnO2  acts  as  a  cata- 
lyser  in  the  reaction  :- 

2H2O2  —  2H2O  +  "Oi 

(d)  Test  for  chromates.     Half  fill  a  test  tube  with  H2O;  add  10  to  15 
drops  of  HC1;  add  a  few  drops  of  potassium  dichromate  solution  to  impart 
a  brick  red  color  to  solution  in  test  tube;  add  a  layer  of  ether  2  cm.  thick; 
add  two  or  three  drops  of  H2O2;  shake.     Is  the  ether  colored  blue  ?     This 
blue  color  is  supposed  to  be  due  to  presence  of  perchromic  acid  (H2Cr2O8), 
probably  formed  as  indicated  by  following  equations: — 

K2Cr2O7  +  2HC1  — .  H2Cr2O7  +  2KC1  (?) 
H2Cr207  +  H202  ->  H2Cr208  +  H2O. 

A  chromate  made  strongly  acid  with  HC1  may  be  used  instead  of  the 
K2Cr2O7  solution.  The  H2O2  acts  as  an  oxidizer. 

(e)  H2O2  as  a  reducing  agent  and  a  bleacher.     To  a  few  cm.3  of  a  potas- 
sium permanganate  add  an  equal  volume  of  H2SO4;  now  add  sufficient 
H2O2  to  decolorize  the  solution. 


HYDROGEN.  7 1 

KMnO4  +  H2SO4  —  HMnO4  +  KHSO4. 
2HMnO4  +  2H2SO4  +  sH2O2  —  2MnSO4  +  8H2O  -f  5O2. 

(/)  Bleaching  with  H2O2,  a  process  of  oxidation.  To  a  few  drops  of 
indigo  solution  add  5  cm.3  of  H2O,  then  add  4  or  5  cm.3  of  H2O2;  heat. 
Results  ? 

C16H10N202  +  2H202-+  2C8H5N02  +  2H2O. 
Experiment  XI. — Preparation  of  Hydrogen  Dioxide. 

Half  fill  a  test  tube  with  H2O;  add  about  20  drops  of  dilute  HC1  and  a 
sufficient  quantity  of  a  potassium  dichromate  (K2Cr2O7)  to  impart  a  brick- 
red  color  to  the  solution;  add  ether  until  there  is  a  layer  2  cm.  thick;  make 
a  thick  paste  of  BaO2  and  H2O;  add  this  paste  to  the  contents  of  the  t.t. ; 
shake  and  observe  the  color  of  the  layer  of  ether.  Compare  results  with 
Exp.  X  (d).  Write  the  equation  for  the  interaction  of  BaO2  and  HC1. 
Another  method  for  the  preparation  of  H2O2  is  represented  by  following 
equation:  Decant  the  solution  which  is  formed  as  the  result  of  the 
reaction, 

Ba02  +  H2S04  dil.  —  Ba2S04  +  H2O2, 

into  a  clean  test  tube;  add  ether,  and  a  few  drops  of  a  dilute  solution  of 
potassium  dichromate. 

H2  +  O2  — *  H2O2,  Aq  +  45,300  cal. 
H2O2  — »  H2O  +O  +  23,100  cal. 
H2  +  O  —  H2O,Aq  +  68,400  cal. 

Experiment  XII. — (Quant.)     Optional.     Synthesis  of  Water. 

Recall  or  repeat  Exp.  "  Combining  Volumes,  "  under  head  of  "  Quanti- 
tative Relationships." 

Experiment  XIII. — (Quant.)  Hydrogen  Equivalent  of  Zinc,  or  a 
Determination  of  the  Volume  of  Hydrogen  Evolved  from  Sulphuric  Acid 
by  a  Given  Weight  of  Zinc. 

Assemble  gas  measuring  apparatus  described  in  Exp.  X.  under  "  Oxy- 
gen." Raise  leveling  bulb  until  water  in  gas  burette  stands  at  top  of 
capillary  on  burette;  clean  and  dry  a  150  cm.3  bottle  which  has  a  small 
test  tube  fused  into  the  bottom  (this  test  tube  need  not  be  fused  to  bottle); 
fit  to  bottle  a  rubber  cork  perforated  with  one  hole;  by  means  of  a  short 
piece  of  capillary  glass  tubing  connect  bottle  with  top  of  burette.  Clean 
and  dry  a  piece  of  pure  sheet  zinc;  weigh  a  piece  of  this  zinc  with  great 
accuracy;  the  piece  must  not  weigh  more  than  .13  grm.  if  the  capacity  of 
burette  is  50  cm.3;  carefully  wrap  a  piece  of  platinum  wire  around  the  zinc 
(the  platinum  acts  as  a  catalyser);  the  piece  of  zinc  should  how  be  dropped 
into  the  bottle;  25  c.m3  of  pure  H2SO4  (i  of  acid  to  4  of  H2O)  is  carefully 
placed  in  the  test  tube  in  the  bottle  by  means  of  a  pipette ;  the  cork  is  forced 
into  place  and  air-tight  connections  are  made  with  the  burette.  It  is  well 
to  place  bottle  in  a  water  bath;  allow  the  apparatus  as  assembled  to  stand 


72  EXPERIMENTAL  CHEMISTRY. 

for  five  minutes;  level;  take  temperatures  of  water  bath  and  water  jacket, 
read  burette,  read  barometer;  record  all  readings;  turn  bottle  until  acid  is 
poured  out  of  t.t.  upon  zinc;  gradually  lower  the  leveling  bulb  as  the  gas 
is  evolved;  after  all  the  zinc  has  been  exhausted  and  gas  is  no  longer 
evolved,  allow  the  apparatns  to  stand  for  a  few  minutes;  bring  the 
water  bath  and  the  water  jacket  to  their  respective  initial  temperatures; 
bring  water  in  burette  and  bulb  to  same  level;  read  burette;  reduce  vol- 
ume of  gas  to  o°  C.  and  760  mm.  If  the  weight  of  i  cm. 3  of  hydrogen 
at  o°  C.  and  760  mm.  is  .00009  grm.  how  many  grams  of  hydrogen  will  be 
evolved  by  32.7  grams  of  zinc? 

Calculate  both  the  weight,  and  the  volume  (at  standard  conditions)  of 
hydrogen  which  should  have  been  evolved  from  the  acid  by  the  weight  of 
zinc  used.  Compare  experimental  data  with  calculated. 

Note. — A  few  drops  of  PtCl4  added  to  the  acid  could  have  been  sub- 
stituted for  the  platinum  wire.  CuSO4  is  also  frequently  used,  but  a 
correction  factor  is  necessary  owing  to  the  following  indicated  reaction — 

Cu  SO4  +  Zn  —  Cu  +  ZnSO4. 

OUTLINE    OF    HYDROGEN. 
I. — History  of  Hydrogen. 

(a)  Discovery — when  and  by  whom? 

(b)  Derivation   of   name.     (Gr.  hudor  (water),  and.  geinomai 

(I  produce).) 

(c)  Historical  in  connection  with  (a)  and  (b). 
II. — Occurrence  or  Distribution. 

III. — Methods  of  Preparation. 
IV. — Physical  Properties. 

Colorless,  Sp.   Ht.   (gas    at  const,   pres.)  3.409. 

Odorless,  Sol'ty  in  Aq.,  1.82  vols.  in  160  (20°  C.). 

Tasteless,  Diffuses  rapidly. 

Wt.  of  il.,  .09  grm.  Melting  point,  — 256°  to  — 257°  C. 

Density  (O  =  1 6),  1.008.  Boiling  point,  — 252°  to  — 253  °C. 

Density  (Air  =  i ),  .0696.  Crit.  Temp.,  —238°  to  —  240°  C. 

Note. — K.  Olszewski,  in  an  attempt  to  liquefy  helium,  cooled  the 
gas  to  — 259°  C.  under  180  atmospheres'  pressure,  by  the  aid  of  solid  hy- 
drogen; the  pressure  was  suddenly  reduced  to  that  of  the  atmosphere 
which  should  give  a  degree  of  cold  as  calculated  by  Laplace  and  Poisson's 
formula,  equal  to  — 271.3°  C.  Helium  did  not  liquefy.  This  is  the 
lowest  temperature  recorded  to  date. 

V. — Chemical  Properties. — Hydrogen,  in  the  majority  of  its  chemical 
relationships,  displays  the  characteristics  of  a  metal,  and  because  of  this 
it  is  usually  regarded  as  a  gaseous  metal*  at  ordinary  temperatures. 

*  NOTE. — Since  hydrogen  does  not  form  a  base  with  oxygen  and  because  its  combi- 
nations with  many  of  the  non-metals  are  acids  and  not  salts,  it  does  not  come  within 
the  ordinary  definition  of  a  metal. 


HYDROGEN.  73 

It  is  practically  diametrically  opposed  to  oxygen.  It  is  liberated  at  the 
negative  electrode  by  the  electric  current  during  electrolysis;  it  is  dis- 
placed from  many  of  its  compounds  by  metals,  i.e.,  hydrogen  functions 
as  a  metal;  sodium  and  potassium  absorb  hydrogen  when  heated  from 
250°  to  400°  C.,  forming  alloys  (Na^H  and  K2H — Richter);  similarly  the 
compound  PdH2  conducts  itself  like  an  alloy  of  two  metals;  according 
to  Graham  the  specific  gravity  of  the  condensed  hydrogen  in  these  com- 
pounds is  found  to  be  .62,  which  makes  it  somewhat  heavier  than  the  metal 
lithium.  Hydrogen  combines  energetically  with  oxygen,  fluorine,  chlor- 
ine and  the  metal  lithium;  however,  it  unites  directly  with  but  few  of  the 
elementary  substances.  Its  " affinity"  for  oxygen  and  chlorine  under 
certain  conditions  is  such  that  it  will  displace  the  elements  with  which 
they  are  united.  When  oxygen  is  taken  away  from  a  compound  by  hydro- 
gen, the  latter  is  said  to  be  oxidized  and  the  compound  reduced. 

The  thermochemical  deportment  of  hydrogen  and  oxygen  in  the  forma- 
tion of  the  two  compounds,  H2O  and  H2O2,  is  interesting. 

H2  +  O  —  H2O  (18  grm.  at  20°  C.)  +  68,360  cals. 
(H2  -f  O)  3  vols.  — >  H2O  (2  vols.  steam)  +  193  cals. 
H2O  (steam,  100°  C.)  —  H2O  (water,  100°  C.)  +  9,666  cals. 
H2O  (water,  100°  C.)  —  H2O  (water,    20°  C.)  +  1,440  cals. 
193  cals.  +  9,666  cals.  +  1,440  cals.  =  11,299  cals- 
68,360  cals.  — 11,299  cals-  =  57>°6i  cals. 

The  approximate  thermal  equivalent  of  the  chemical  energy  of  a  mix- 
ture of  2  grm.  of  hydrogen  and  16  grm.  of  oxygen  is  57,061  cals. 

The  "Law  of  Hess"  is  suggested  by  the  thermo-chemical  conduct  of 
hydrogen  dioxide: 

(i)  H2  +  02  —  H202Aq  +  45,300  cal. 
and     (2)  H2O2  —  H2O  +  O  +  23,100  cal. 


adding  (3)  H2  +  O    — »    H2O  +  68,400  cal. 

No  matter  how  many  stages  there  are  to  a  given  reaction,  if  the  initial 
and  final  states  are  the- same  in  each  case,  the  heat  of  the  reaction  will 
be  constant. 

On  inspecting  equation  (2)  it  is  obvious  that  H2O2  is  an  active  oxidizer, 
as  it  liberates  23,100  cal.  more  than  when  the  same  amount  of  free  oxy- 
gen is  used  under  similar  conditions.  This  serves  to  explain  its  activity 
as  an  oxidizer  in  foregoing  experiments.  Which  is  the  more  active  oxid- 
izer O3  or  H2O2  ?  Hint — compare  energy  equations. 

Compounds  of  Hydrogen. — To  be  studied  as  work  progresses. 

Uses  of  Hydrogen. — :Oxy-hydrogen  blowpipe,  calcium  light,  etc. 

Miscellaneous  Topics. — Principles,  theories,  definitions,  etc. 

Note. — In  the  future  the  student  will  make  a  brief  written  " resume" 
of  the  study  of  each  element.  It  should  embrace  the  various  heads,  under 
which  data  is  to  be  tabulated,  as  suggested  above  in  the  "  Outline  of 
Hydrogen." 


74  EXPERIMENTAL  CHEMISTRY. 

PROBLEMS. 

i. — What  is  the  weight  of  100  cm.3  of  hydrogen  at  20°  C.  and 
760  mm.  ? 

2. — How  many  grams  of  zinc  will  be  required  to  liberate  10  grams 
of  hydrogen  from  sulphuric  acid? 

3. — What  is  the  weight  of  a  liter  of  hydrogen  measured  over  water  at 
20°  C.  and  777.36  mm.? 

4. — How  many  liters  of  hydrogen  can  be  obtained  from  100  cm.3  of 
sulphuric  acid  (density,  1.84)? 

5. — How  much  zinc  and  sulphuric  acid  will  be  required  (theoretically) 
to  liberate  i  liter  of  hydrogen? 

6. — A  liter  of  oxygen  weighs  as  much  as  what  number  of  liters  of 
hydrogen  ? 

7. — Two  grams  of  hydrogen  are  equivalent  to  how  many  liters  of  the 
gas  at  standard  conditions? 

8. — What  are  the  valences  of  the  elements  in  the  following:  HC1, 
H2O,  NasO,  NH3,  KC1,  LiH,  H2S,  CO2,  FeO,  Fe^? 

9. — If  a  liter  of  hydrogen  weighs  .09  grm.  what  is  the  weight  of  a  liter 
OfO2?  OfCO2?  OfN2?  OfCLj? 

10. — What  are  the  relative  rates  of  effusion  of  hydrogen,  oxygen  and 
carbon  dioxide? 

ii. — Why  is  the  oxygen  admitted  through  the  inner  tube  of  the  oxy- 
hydrogen  blowpipe?  Note:  The  temperature  of  the  hydrogen  flame 
in  air  is  about  2000°  C.;  in  oxygen  it  is  about  2500°  C.  It  can  not  sur- 
pass this  latter  temperature  as  it  is  the  temperature  at  which  steam  is  re- 
solved into  its  elements. 

12. — What  are  the  formulae  of  cuprous  and  cupric  oxides?  Of  fer- 
rous and  ferric  oxides?  Of  ferrous  and  ferric  chlorides? 

13. — Is  the  process  of  oxidation  accompanied  by  the  process  of  reduc- 
tion ?  Illustrate  by  equations  and  interpret. 

14. — Is  occlusion  a  chemical  or  physical  action  ?     State  your  reasons. 

15. — Mention  a  reaction  in  which  the  speed  was  altered  by  surface 
effects. 


CHAPTER  X. 
WATER. 

In  all  investigations,  the  product  of  the  interaction  of  two  volumes  of 
hydrogen  and  one  volume  of  oxygen,  has  been  proven  to  be  identical 
with  the  substance  which  is  known  as  water.  It  is  one  of  the  most 
abundant  and  universally  distributed  of  all  chemical  compounds.  It  is 
essential  to  life.  Its  properties  are  remarkable  and  diversified. 

In  1781,  Cavendish  confirmed  the  formation  of  water  by  the  com- 
bustion of  hydrogen.  Prior  to  his  work,  water  was  thought  to  be  an  ele- 
mentary substance.  Lavoisier  determined  its  quantitative  composition 
in  1783.  In  1805,  Gay-Lussac  showed  that  it  was  produced  by  the  union 
of  two  volumes  of  hydrogen  with  one  volume  of  oxygen.  Water  was  first 
decomposed  by  electricity  in  1800  by  Nicholson  and  Carlisle.  Davy  con- 
firmed and  extended  the  work  of  these  men,  aside  from  initiating  many 
brilliant  experiments.  The  names  of  the  men  who  were  connected  with 
the  study  of  water  are,  Cavendish,  Priestly,  Lavoisier,  Humbolt,  Gay- 
Lussac,  Nicholson  and  Carlisle,  Berzelius,  Davy,  Dumas,  Dulong,  Stas, 
and  Morley. 

Experiment  I. — Composition  of  Water. 

(a)  Synthesis  of  water.     Recall  or  repeat  the  various  experiments  in- 
volving the  " synthesis  of  water." 

2H2  (two  vols.)  +  O2  (one  vol.)  — *  2H2O  (two  vols.) 
Recall  "  Law  of  Combining  Volumes."     State  it. 

(b)  Analysis  of  water. 

1.  By  electrolysis. 

2.  By  action  of  metals. 

4H20  (steam)  +  3  Fe  (hot)  <^Fe3O4  +  4H2  (Fig.  18). 

At  a  high  temperature,  steam  oxidizes  those  elements  which  readily 
combine  with  oxygen.  Recall  the  respective  oxidizing  activities  of  O2, 
O3  and  H2O2.  Compare  their  respective  energy  equations  with  that  of 
steam  when  each  is  acting  as  an  oxidizing  agent. 

Experiment  II. — Physical  Properties. 

(a)  Quant.)  Density.  Clean  and  dry  a  50  cm.3  Erlenmeyer  flask; 
weigh  it.  By  means  of -a  burette  or  a  pipette  introduce  into  the  flask  15 
cm. 3  of  distilled  H2O  which  has  a  temperature  of  about  20°  C.;  weigh 
flask  and  contents  as  rapidly  as  possible  to  prevent  loss  by  evaporation. 
Calculate  the  approximate  density  of  H2O  under  the  existing  conditions. 
See  "Table  of  Density"  in  Appendix. 

75 


70  EXPERIMENTAL  CHEMISTRY. 

Note. — Burettes  filled  with   distilled  H2O  may  be  placed  in  readily 
accessible  places  in  the  laboratory. 

(b)  Freezing  point.     Fill  a  beaker  with  a  mixture  of  clean  ice  and  water; 
the  ice  should  be  broken  into  small  pieces.     Suspend  a  thermometer  in 
the  mixture  for  5  or  10  min.;  the  mixture  should  be  stirred.     Tap  the 
thermometer  with  finger  and  read.     Record  reading. 

(c)  Boiling  point.     Measure  20-25  cm-3  °f  distilled  H2O  into  a  clean 
Erlenmeyer  flask;  suspend  a  thermometer  in  the  flask  so  that  its  bulb  is 
about  2  or  3  cm.  above  surface  of  H2O;  heat  flask  and  contents  until  there 


FIG.  18.— (Smith  and  Keller.) 


is  a  rapid  evolution  of  steam;  tap  thermometer,  and  read.  Record 
reading.  Suspend  thermometer  so  that  bulb  dips  into  water  but  does 
not  touch  bottom  of  flask.  Repeat  above.  Record  reading.  Read 
the  barometer.  (Instructions.)  Record  reading. 

Experiment  III. — Purity  of  Water. 

Place  a  few  drops  of  distilled  water  upon  a  clean  "watch  glass"  and 
evaporate  to  dryness  upon  a  steam  bath.  Is  there  a  stain  or  residue  on 
glass  ?  Repeat  above  using  ordinary  water,  for  example,  "  drinking  water." 
Compare  results  with  above.  If  the  impurities  of  the  water  are  very  vola- 
tile will  this  method  enable  you  to  detect  them  ?  Why  ? 

The  preparation  of  "absolutely"  pure  water  is  an  impossibility,  as  the 
material  of  any  vessel  is  soluble  in  a  greater  or  lesser  degree. 

Experiment  IV.— "Temperature -Density  "  Graph  of  Water. 

By  referring  to  the  table  in  the  Appendix,  plot  on  "coordinate  paper" 
the  "temperature-density"  graph  of  water.  Use  the  axis  of  ordinates  for 
the  scale  of  density  and  the  axis  of  abscissas  for  temperature-scale. 
(Instructions.) 

Does  the  comparatively  abrupt  change  in  the  direction  of  graph  sug- 
gest the  introduction  of  a  new  "factor"  in  the  phenomena  which  graph  is 
supposed  to  represent?  May  this  change  in  direction  merely  represent 
that  a  factor  or  a  group  of  factors  has  suddenly  become  dominant? 


WATER.  77 

Experiment  V. — Purification  of  Water. 
I. — Chemical  Methods. 
II. — Mechanical  Methods. 
Decanting. 
Filtering. 
Boiling. 
Distillation. 

(a)  Distillation  of  a  solution  of  copper  sulphate.  Assemble  a  distilling 
apparatus,  composed  of  a  distilling  flask,  Liebig  condenser  (Fig.  19)  and 
a  receiver,  or  use  a  retort  and  a  receiving  flask.  (Instructions.)  Test 
solution  with  litmus  paper;  test  distillate  with  litmus  paper;  compare 
results.  How  does  color  of  distillate  compare  with  color  of  original 
solution  ? 


FIG.  19. 


(b)  Distillation  of  a  solution  of  ammonium  hydroxide.     Proceed  as  in 
(a)  Results?     Explain. 

(c)  Distillation  of  a  solution  of  alcohol.     Proceed  as  in  (a),  but  test  in- 
flammability of  solution  and  distillate  by  means  of  a  lighted  match. 
Explain.     Recall  "Law  of  Partial  Pressures." 

Note. — Boiling-tubes  prevent  "bumping"  of  solutions  undergoing 
ebullition. 

Experiment  VI. — Hydrolysis. 

To  a  small  quantity  of  bismuth  chloride,  Bi  C13,  or  antimony  chloride, 
SbCl3,  add  5  cm.3  of  water;  the  white  precipitate  is  essentially  bismuth 
oxychloride,  BiOCl.  To  the  precipitate  add  a  few  drops  of  concen- 
trated hydrochloric  acid,  then  warm  tube;  by  repeating  this  process 


78  EXPERIMENTAL  CHEMISTRY. 

dissolve  the  precipitate  in  a  minimum  quantity  of  acid.  The  tube  con- 
tains a  solution  of  SbCl3.  To  a  test  tube  nearly  full  of  water  add  a 
few  dr,ops  of  the  solution.  Explain  the  formation  of  the  white  precipitate. 
Write  equations  for  the  three  reactions. 

Experiment  VII. — Chemical  Union  of  Water  with  Oxides. 

(a)  Combination  with  a  metallic  oxide.     Place  a  small  quantity  of 
quick  lime  (CaO)  or  barium  oxide  (BaO)  in  a  test  tube;  add  5-10  cm.3 
of  H2O  which  has  been  previously  tested  with  litmus  paper;  shake  vigor- 
ously; let  contents  settle;  decant  clear  liquid  into  another  test  tube  and 
test  liquid  with  litmus  paper.     Results  ?     Equation  ?     Substances  like 
the  above  which  turn  red  litmus  blue  are  called  "bases.'' 

(b)  Combination  with  a  non-metallic  oxide.     Recall  the  action  of  P2O5 
and  SO2  on  the  moist  litmus  paper.     Record  results.     Substances  formed 
by  the  union  of  non-metallic  oxides  and  water  and  which  turn  blue 
litmus  red,  are  called  "acids." 

Experiment  VIII. — Water  of  Hydration  (Crystallization). 

(a)  Place  a  small  crystal  of  copper  sulphate  (CuSO4.5H2O)  in  a  test 
tube;  heat  gently  until  a  white  powder  remains.     Is  there  evidence  that 
water  has  been  liberated  during  the  above   process  ?     Add  a  drop  or  two  of 
water  to  the  powder  when  tube  is  cool.     Effect  ?     Add  water  and  boil,  dis- 
solving the  powder  in  the  least  possible  quantity  of  water;  set  tube  aside 
for  several  days;  do  crystals  form? 

(b)  Using  crystals  of  gypsum,  potassium  dichromate,  barium  chloride 
and  potassium  nitrate,  ascertain  whether  water  is  present  in  each  crystal. 

Note. — Some  crystals  contain  " mechanically  inclosed"  water;  when 
such  crystals  are  heated  they  fly  to  pieces  explosively:  they  are  said  to 
decrepitate.  Did  any  of  the  above  crystals  decrepitate? 

Do  all  crystals  contain  water  oj  hydration  ?     Reasons  for  your  answer  ? 

A  salt  containing  water  of  hydration  is  spoken  of  as  a  hydrated  salt; 
when  the  water  has  been  removed  it  is  known  as  an  anhydrous  or  dehy- 
drated salt. 

Experiment  IX. — Vapor  Tension  of  Substances.  Efflorescence  and 
Deliquescence. 

(a)  Examine  a  small  clear  crystal  of  sodium  sulphate,  Na.jSO4.io  H2O; 
place  it  on  a  clean  watch  glass  and  set  it  in  closet;  after  several  days  ex- 
amine.    Explain.     Equation  ?     The  substance  is  said  to  be  efflorescent. 

(b)  Place  a  piece  of  dehydrated  calcium  chloride,  CaCl2  in  a  beaker. 
Repeat     (a).     Explain.     Equations?     The    substance    is    said    to    be 
deliquescent. 

(c)  Introduce  10  cm.3  of  concentrated  sulphuric  acid  into  a  dry  test 
tube;  mark  its  height  by  means  of  a  piece  of  label  or  an  ink  mark;  allow 
it  to  stand  for  a  week.     Explain. 


WATER.  79 

Experiment  X.—  (Quant.)     Determination  of  Water  of  Crystallization. 

Weigh  accurately  a  clean  dry  crucible  and  cover.  Record  weight. 
Introduce  into  crucible  about  2  grm.  of  powdered  copper  sulphate  crys- 
tals ;  place  cover  on  crucible  and  weigh  accurately.  Record  weight.  Place 
covered  crucible  on  a  pipe-stem  triangle  and  heat  gently  for  about  25  min. 
or  until  quite  certain  that  the  blue  color  has  completely  disappeared; 
cool  in  desiccator,  then  weigh.  Record  weight.  Heat  again  for  5-10 
min.  and  weigh.  Repeat  until  weight  becomes  constant.  Calculate 
the  percentage  of  water  of  crystallization  of  the  crystals. 

Note.  —  In  applying  heat  to  the  crucible,  the  tip  of  the  flame  may 
barely  touch  the  bottom  of  the  crucible. 

Physical  Properties  of  Water. 

Tasteless.  Ht.  of  Fusion,  80  cals. 

Odorless.  Ht.  of  Vaporization,  536.7  cals. 

Colorless  (in  thin  layers).  Melting  point,  o°  C.  760  mm. 
Density  (i  cm.3  at  4°  C.),  i  grm.  Boiling  point,  100°  C.  760  mm. 

Sp.  Ht.  (Solid  state),  .50.  Crit.  temp.,  370°  C.  (Highest  critical 
Sp.  Ht.  (Liquid),  i.oo  temperature  known.) 

Sp.  Ht.  (Gas),  .477- 

The  specific  heat  of  water  is  remarkably  high.  More  heat  is  required 
to  raise  the  temperature  of  a  given  weight  of  it  one  degree  than  is  re- 
quired for  any  other  substance  except  hydrogen. 

Chemical  Properties.  —  Water  is  one  of  the  most  stable  of  all  substances, 
i.e.,  it  is  not  easily  decomposed  as  regarded  from  the  standpoint  of  energy, 
yet  like  other  chemical  compounds  it  is  broken  up  into  its  elements  by 
heat.  "Sainte-  Claire  Deville  was  the  first  to  carefully  investigate  and 
explain  the  decomposition  of  water  by  pouring  molten  platinum  (1770°  C.) 
into  it."  "  He  proved  that  the  dissociation  —  a  reversible  decomposition  — 
did  not  take  place  suddenly,  but  gradually;  that  it  advanced  regularly 
with  increasing  temperature,  and  was  limited  by  an  opposing  combina- 
tion-tendency on  the  part  of  the  components."  —  Richter. 


The  decomposition  is  appreciably  initiated  at  1000°  C.  and  is  about  half 
complete  at  2500°  C.  The  percentage  of  dissociation  increases  with  in- 
crease of  temperature. 

Water  bears  evidence  of  being  slightly  dissociated  when  in  the  liquid 
condition  as  well  as  when  in  the  gaseous  state.  Electrical  conductivity 
and  certain  chemical  reactions  indicate  the  truth  of  the  following  equation: 

H2O<=±H+  +  OH'. 

This  will  be  discussed  more  fully  under  the  subject  of  "The  Modern 
Theory  of  Solution."  Water  is  usually  referred  to  as  a  perfectly  neutral 
substance,  but  this  is  far  from  agreeing  with  experimental  facts.  Owing 


8o  EXPERIMENTAL  CHEMISTRY. 

largely  to  the  products  of  dissociation,  water  manifests  definite  chemical 
properties  under  definite  conditions. 

Another  interesting  property  of  water  is  one  which  is  thought  to  be 
closely  related  to  the  varying  density  of  water  with  change  of  temperature. 
In  virtue  of  this  property  the  molecules  of  water  are  supposed  to  associate, 
that  is,  a  number  of  molecules  combine  and  form  a  "  molecular-complex." 
This  phenomenon  is  known  as  the  polymerization  of  water.  "  Surface 
tension"  and  "depression  of  the  freezing  point"  experiments  support  this 
view.  Raoult  and  others  hold  that  at  temperatures  near  the  freezing 
point,  a  relatively  large  quantity  of  the  water  possesses  the  molecular 
formula,  H8O4  or  (H2O)4.  Increase  in  temperature  causes  a  breaking 
up  of  these  molecular-complexes,  and  vice  versa.  The  change  in  density 
of  water  with  change  of  temperature  is  explained  in  general  as  follows: 

The  H8O4  molecules  occupy  more  space  than  4H2O;  polymerization 
increases  as  temperature  is  lowered;  water  contracts  when  the  temperature 
falls;  two  factors  must  then  be  considered.  Water  expands  when  cooled 
below  4°  C.  because  the  effects  of  polymerization  overbalance  the  effects 
due  to  contraction  as  the  result  of  lowering  the  temperature.  When  the 
temperature  is  increased  to  4°  C.,  the  water  should  expand  normally, 
and  probably  does,  but  the  molecular-complexes  tend  to  split  up  which 
per  se  causes  a  shrinkage  in  volume;  at  4°  C.  the  effects  of  the  two  factors 
balance  each  other,  hence  the  maximum  density  at  this  temperature.  If 
the  temperature  is  raised  above  4°  C.,  then  the  expansion-effects  due  to  in- 
crease of  temperature  are  greater  than  the  contraction-effects  due  to  a 
breaking  down  of  the  associated  molecules.  Above  4°  C.  the  effect  of 
cubical  expansion  due  to  increased  temperature  is  dominant;  below  4°  C., 
expansion-effects  due  to  increased  polymerization  are  dominant. 

It  was  said  in  the  beginning  of  this  chapter  that  water  in  its  chemical 
relations  presents  some  very  remarkable  features;  one  of  the  most  marked 
is,  that  although  it  is  an  indifferent  oxide  (hydrogen  oxide),  it  possesses 
a  group  of  combining  tendencies  which  extend  over  a  wider  range  than 
those  of  any  other  chemical  compound.  It  combines  directly  with  many 
compound  substances,  but  with  few  elements — chlorine  and  bromine  being 
perhaps  the  only  elementary  substances.  Gases  are  dissolved  by  water 
to  a  relatively  slight  extent,  but  it  can  not  be  affirmed  positively  that  they 
enter  into  a  chemical  combination.  The  most  common  of  its  chemical 
reactions  is  the  formation  of  a  class  of  compounds  known  as  hydrates, 
which  exist  in  the  solid  form  only,  undergoing  decomposition  when  placed 
in  solution,  Ex.s.  CuSO4.5H2O,  Na^SO^io  H2O.  They  show  a  definite 
composition  and  frequently  much  heat  is  developed  during  their  for- 
mation, thus  (Na^COg,  10  H2O),  equals  8800  cal.  Water  also  reacts  with 
some  substances  in  a  manner  which  is  very  similar  to  metathesis — to  say 
that  the  products  of  dissociated  water,  rather  than  the  water  per  se,  inter- 
acts with  substances  to  produce  double  decomposition,  is  a  statement  more 
nearly  in  accord  with  facts.  This  kind  of  a  reaction  is  known  as  hydrol- 
ysis. The  mechanism  of  the  interaction  may  be  interpreted  more  readily 
by  the  aid  of  an  equation. 


WATER.  8 1 

H20  <=±  H+  +  OH 

3d'  +  2H+  +  2QH/<=>Bi(OH2)Cl  +  2H+'  +  2Cl' 
Bi(OH)2Cl<=±BiOCL  +  H2O. 

If  the  product  of  hydrolysis  is  soluble  there  is  usually  little  or  no  visible 
evidence  of  the  interaction;  if  the  product  is  insoluble  a  precipitate  forms. 
As  the  work  advances  the  phenomenon  of  hydrolysis  will  be  more  thor- 
ough discussed.  A  familiar  example  of  the  direct  combination  of  water 
and  metallic  and  non-metallic  oxides  is  the  slaking  of  quicklime  which 
may  be  represented  as  follows: 

CaO+  H2O  —  Ca(OH)2  (Calcium  Hydroxide), 
again,     SO2  +  H2O  —  H2SO3  (Sulphurous  Acid). 

The  product  of  the  first  reaction  belongs  to  a  group  of  substances  known 
as  bases;  the  latter  product,  to  a  group  whose  generic  name  is  acids. 
The  chemical  properties  of  these  two  groups,  acids  and  bases,  are  very 
different;  an  aqueous  solution  of  the  former  turns  red  litmus  paper  blue, 
while  in  the  case  of  the  acids,  blue  litmus  paper  is  turned  to  a  red  color. 

However,  of  all  the  various  properties  of  water,  perhaps  none  are  of 
more  importance  to  the  chemist  than  its  solvent  properties.  The  question 
as  to  whether  the  effecting  of  a  solution  is  a  chemical  or  physical  process 
has  not  been  satisfactorily  answered.  This  question  will  be  discussed 
more  fully  in  the  next  chapter. 

In  addition  to  its  enumerated  uses,  it  should  be  remembered  that  water 
is  the  standard  of  many  physical  measurements . 


CHAPTER  XI. 
SOLUTIONS. 

We  have  observed  that  many  gases,  liquids  and  solids  when  placed 
in  water  disappear  and  form  homogeneou  systems  which  are  known  as 
solutions.  The  operation  of  preparing  a  solution  is  called  "dissolving" 
or  " putting  into  solution."  The  substance  dissolved  is  known  as  the 
solute  and  the  material  in  which  the  solute  is  dissolved,  the  solvent. 
These  two  terms  are  unfortunate  inasmuch  as  they  do  not  suggest  the 
mutual  interaction  of  solute  and  solvent  during  the  process  of  dissolving. 
All  substances  are  soluble  in  a  degree, — the  solubility  depending  upon  the 
relative  strength  of  the  affinities  oj  the  substance  for  itself  and  for  the  solvent. 
If  the  affinity  of  solvent  and  solute  is  greater  than  the  affinity  of  sub- 
stance for  self,  then  the  solubility  will  be  correspondingly  great.  The 
degree  of  solubility  is  usually  expressed  by  the  terms  insoluble,  slightly 
soluble,  soluble,  and  very  soluble.  When  the  solution  contains  a  rela- 
tively large  quantity  of  the  solute  it  is  said  to  be  a  concentrated  solution; 
if  a  relatively  small  quantity,  a  dilute  solution. 

"  Since  matter  in  every  state  can  be  mixed  with  other  matter,  irre- 
spective of  its  state,  it  is  obvious  that  many  different  kinds  of  solutions 
are  possible."  Jones  (H.  G.)  gives  the  following  list: 

I.  Solution  of  a  solid  in  a  solid. 
II.  Solution  of  a  solid  in  a  liquid. 

III.  Solution  of  a  solid  in  a  gas. 

IV.  Solution  of  a  liquid  in  a  solid. 
V.  Solution  of  a  liquid  in  a  liquid. 

VI.  Solution  of  a  liquid  in  a  gas. 

VII.  Solution  of  a  gas  in  a  solid. 

VIII.  Solution  of  a  gas  in  a  liquid. 

IX.  Solution  of  a  gas  in  a  gas. 

Among  the  properties  of  solutions  may  be  noted — color,  odor,  taste, 
density,  expansibility,  compressibility,  surface  tension,  viscosity,  vapor 
tension,  osmotic  pressure,  refractive  index,  definite  boiling  and  freezing 
points,  etc. 

SOLUBILITY    OF    SOLIDS. 

Experiment  I. — Characteristics  of  a  Solution. 

Powder  separately  a  few  small  crystals  of  alum,  and  potassium  di- 
chromate;  dissolve  each  in  the  least  possible  quantity  of  water— the  test 
tubes  should  be  shaken  and  warmed  repeatedly  to  aid  in  the  process. 
Are  the  solutions  clear?  Transparent?  Homogeneous?  Pour  the 
solutions  into  separate  crystallizing  dishes,  or  set  the  tubes  aside  to  cool. 

82 


SOLUTIONS.  83 

Do  crystals  form  ?     If  so,  decant  the  liquid  and  examine  crystals.     Are 
they  similar  to  the  original  crystals  ? 

A  solution  is  ordinarily  defined  as  a  clear  transparent  homogeneous 
mixture,  the  components  of  which  can  not  be  separated  by  a  purely 
mechanical  process.  "Colloidal"  or  "pseudo-solutions" — see  Ostwald's 
"Principles  of  Chemistry."  to 

Experiment  II. — Surface  and  Diffusion  Phenomena.     Solution  Tension. 

Half  fill  a  test  tube  with  water;  drop  one  small  crystal  of  potassium 
permanganate  into  the  water;  set  tube  where  it  will  not  be  shaken,  yet 
can  be  easily  observed;  as  you  continue  your  experimenting,  notice  the 
color  of  the  solution  at  short  intervals  of  time.  Record  observations. 
Was  a  relatively  long  period  of  time  required  for  the  dissolving  of  the 
crystals  ?  Postulating  that  the  salt  is  quite  soluble  in  water  at  a  given 
temperature,  what  factors  determine  largely  the  speed  of  the  dissolv- 
ing process?  To  what  two  mechanical  processes  do  we  resort 
usually  in  an  endeavor  to  hasten  solution?  Did  any  phenomena  occur 
during  the  dissolving  of  the  solid  which  would  suggest  any  of  the  prop- 
erties possessed  by  liquids  or  gases  ?  Enumerate  them. 

The  process  of  solution  receives  a  partial  explanation,  at  least,  when 
the  accompanying  phenomena  are  interpreted  in  terms  of  the  kinetic- 
molecular  hypothesis.  When  a  soluble  crystalline  substance  is  intro- 
duced into  the  solvent,  it  is  thought  that  the  molecules  of  which  the 
solute  is  composed  are  detached  and  enter  the  solvent.  By  the  process 
of  diffusion  these  detached  particles  move  away  from  the  surface  of  the 
solute,  and  are  scattered  throughout  the  solvent.  After  a  time  some  of 
these  particles,  which  move  in  every  direction,  will  again  come  into  con- 
tact with  the  solid  solute  and  attach  themselves  to  it.  If  the  rate  at  which 
the  molecules  press  into  solution  is  greater  than  that  with  which  they 
return  to  the  solute,  it  is  evident  that  the  solute  will  eventually  be  wholly 
dissolved.  If  a  sufficient  quantity  of  the  solid  substance  has  been  placed 
in  the  solvent  the  dissolving  process  will  continue  until  the  speeds  of  the 
two  opposing  actions  are  identical.  When  this  occurs  the  solution  is 
said  to  be  saturated.  The  solid  solute  is  then  in  equilibrium  with  the 
dissolved  portion, 

KMn04  (Solid)  <=>  KMnO4  (Dissolved). 

If  this  equilibrium  is  disturbed  by  varying  the  concentration  of  the 
solution,  the  speeds  of  the  opposing  actions  will  be  altered  until  equi- 
librium is  again  established,  either  by  the  dissolving  of  more  of  the 
solute  if  the  concentration  is  decreased,  say  by  diluting  the  solution; 
or  by  the  deposition  of  a  portion  of  the  solute  if  the  concentration  is 
increased  by  any  cause  whatever — say  the  removal  of  a  portion  of  the 
solvent.  The  experimenter  is  never  assured  of  the  existence  of  an 
equilibrium  unless  a  portion  of  the  solute  (solid)  is  in  contact  with  the 


84  EXPERIMENTAL  CHEMISTRY. 

solution.  Powdering  the  solute  and  shaking  the  mixture  hastens  the 
process  of  solution.  Why? 

This  tendency  of  the  molecules  of  a  substance  when  placed  in  a  solvent, 
to  leave  the  solid  and  pass  into  solution  is  referred  to  as  solution  tension 
because  of  its  evident  analogy  to  the  vapor  tension  of  liquids,  in  virtue 
of  which  liquids  tend  to  assume  the  gaseous  condition.  The  analogy  is 
continued:  There  is  equilibrium  when  the  vapor  tension  of  a  liquid  is 
balanced  by  the  gaseous  pressure  of  the  vapor  above  it;  likewise  there  is 
a  dynamic  equilibrium  in  a  saturated  (concentrated)  solution  between 
the  dissolved  and  undissolved  portions  of  the  solute;  the  force  (energy) 
in  virtue  of  which  the  molecules  tend  to  pass  into  solution  is  usually 
spoken  of  (mentioned  previously)  as  the  solution  tension  of  the  solute;  the 
force  in  equilibrium  with  the  solution  tension,  and  which  is  considered 
as  the  analogue  of  the  gaseous  pressure  of  the  vapor  above  a  liquid,  is 
known  as  osmotic  pressure.  In  this  pressure  we  recognize  the  cause  of 
the  diffusion  of  substances  in  solution.  It  has  been  discovered  that  the 
molecules  of  a  dissolved  substance  like  sugar,  exert  a  pressure  on  the 
solvent  identical  with  the  pressure  which  they  would  exert  on  the  sides 
of  a  vessel  of  the  same  volume  as  that  of  a  solution,  if  they  were  in  the 
gaseous  state.  This  pressure  is  given  the  name  osmotic  pressure  because 
it  is  only  by  taking  advantage  of  the  phenomenon  of  osmosis  (which 
provides  for  the  elimination  of  "surface  pressures")  that  it  can  be 
rendered  apparent  and  directly  measured. 

The  correlation  of  much  experimental  data  which  will  be  discussed 
as  the  work  proceeds,  has  given  rise  to  what  is  known  as  the  "  physical 
theory"  of  dilute  solutions.  It  may  be  stated  as  follows:  "The  mole- 
cules of  the  dissolved  substance  pervade  the  solvent  without  being 
influenced  thereby,  and  possess  the  same  properties  as  they  would 
possess  did  they  alone,  in  the  state  of  gas  occupy  the  volume  filled  by 
the  solution." — Bloxam.  It  seems  that  there  is  more  likelihood  of  the 
dissolved  substance  being  in  a  condition  comparable  to  a  gas  than  to 
either  a  liquid  or  a  solid.  It  is  obvious  that  it  can  not  be  in  a  state  of 
aggregation  comparable  to  a  solid. 

Hulett  (J.  Amer.  Chem.  Soc.,  27,  49,  1905)  recently  presented  an 
interesting  article  on  the  importance  of  the  state  of  the  solid  from  which 
the  solution  is  made.  The  real  subject  of  the  article,  however,  is  an 
explanation  of  the  greater  speed  of  solution  and  greater  solubility  of  very 
small  particles  of  a  solid  as  "based  on  the  following  considerations: 
The  boundary  between  a  solid  and  a  liquid  is  the  seat  of  a  certain  amount 
of  energy  due  to  the  surface-energy  of  the  liquid;  if  this  surface  is  increased 
by  powdering  the  solid,  the  total  surface-energy  is  correspondingly  in- 
creased. Further,  it  is  a  generally  observed  fact  that  the  form  of  a  sub- 
stance which  has  the  greater  free  energy  is  the  more  soluble,  has  the 
greater  vapor-pressure,  and  is  the  least  stable  form,  e.g.,  allotropic 
modifications  of  substances  have  different  solubilities,  and  the  unstable 
form  is  always  the  more  soluble.  This  phenomenon  js  hardly  analogous 
to  the  well  known  behavior  of  liquid  drops  of  different  sizes.  Small 


SOLUTIONS.  85 

drops  in  the  vicinity  of  large  ones  grow  small  and  disappear,  while  the 
larger  ones  grow  larger,  and  the  reason  is  quite  clear.  It  is  known  that 
the  curved  surface  of  a  liquid  has  a  greater  vapor-pressure  than  a  plane 
or  less  curved  surface;  therefore,  a  distillation  takes  place.  The  simi- 
larity between  the  vapor-pressure  of  liquids  and  the  solution-pressure 
of  solids  has  suggested  to  some  the  analogy  between  the  facts  just  men- 
tioned and  the  behavior  of  solid  particles  of  different  sizes  in  contact 
with  the  solution.  But  we  can  not  assume  that  the  surface  of  the  par- 
ticles of  a  powder  is  curved,  or,  if  that  is  granted,  we  do  not  know  that  a 
curved  surface  of  a  solid  or  a  sharp  edge  has  a  greater  solution-pressure 
than  a  plane  surface  of  the  same  substance. 

"Hulett  found  that  a  solution  of  gypsum  saturated  at  25°,  containing 
2.080  grm.  CaSO4  per  liter  will  increase  its  concentration  rapidly  to  a 
maximum  when  shaken  with  powdered  gypsum  and  then  will  decrease 
to  the  original  value  again.  In  one  experiment  the  content  of  gypsum 
reached  2.542  grm.  CaSO4  in  a  liter  in  a  minute.  The  fine  powder  used 
for.  this  purpose  was  found,  after  the  concentration  had  reached  its 
original  point,  to  have  increased  in  size  of  its  grain.  The  smallest 
particles  thus  go  into  solution,  produce  the  supersaturation  and  are 
then  deposited  upon  the  others,  so  that  all  increase  in  size." — Morgan. 
Hulett  suggests  the  preparation  of  saturated  solutions  from  large  particles 
as  there  is  less  likelihood  of  supersaturation. 

OSMOSIS. 

"  If  a  solution  and  the  pure  solvent  are  separated  by  a  semipermeable 
membrane  the  solvent  will  flow  through  the  membrane  into  the  solution, 
where  its  escaping  tendency  is  less.  The  only  way  of  preventing  this 
flow  is  to  make  the  escaping  tendency  of  the  solvent  the  same  on  both 
sides  of  the  membrane.  There  are  two  simple  ways  of  accomplishing 
this,  (i)  to  increase  the  pressure  on  the  solution  until  the  escaping 
tendency  of  the  solvent  in  the  solution  is  raised  to  equal  that  of  the 
solvent  in  the  pure  state,  (2)  to  diminish  the  pressure  on  the  pure  solvent 
until  its  escaping  tendency  is  lowered  to  equal  that  of  the  solvent  in  the 
solution. 

The  osmotic  pressure  may,  therefore,  be  defined  in  two  ways,  (i)  as 
usually  defined,  it  is  the  increase  in  the  pressure  on  the  solution  necessary 
to  bring  the  latter  into  equilibrium  with  the  solvent;  (2)  Noyes,*  however, 
prefers  to  define  the  osmotic  pressure  as  the  diminution  in  the  pressure 
on  the  solvent  necessary  to  bring  it  into  equilibrium  with  the  solution. 
Neither  of  these  definitions  is  entirely  free  from  objections." — Lewis,  G.  N.f 

Experiment  III. — (L.  T.)  Osmotic  Pressure. 

Note. — The  instructor  or  assistant  should  assemble  the  apparatus 
(Fig.  20)  in  the  presence  of  the  members  of  the  class.  The  use  of  a  solu- 

*Z.  physik.  Chem.,  35,  707  (1900). 

t  Osmotic  Pressure  of  Concentrated  Solutions." — Jour.  Amer.  Ch3m.  Soc.,  30,  668 
(1908). 


86 


EXPERIMENTAL  CHEMISTRY. 


tion  of  sugar  is  suggested.  Whether  the  experiment  is  designed  to  show 
relationships  in  a  quantitative  or  qualitative  manner,  the  effects  of  concen- 
tration and  temperature  upon  the  osmotic  pressure  should  be  demon- 
strated. These  effects  should  be  compared  with  corresponding  effects 


FIG.  20. 

produced  on  gases  by  similar  influences.  By  use  of  the  following  data 
the  possibility  of  applying  the  "gas  law,"  PV  =  RT,  to  dilute  solutions, 
may  be  shown. 

OSMOTIC  PRESSURE  OF  CANE  SUGAR. 


("  Osmotic  Investigations  "- 

Ejject  oj  Concentration. 

Concentration  in  Pressure  in 

per  cent,  by  weight.  mm.  of  Hg. 

1  per  cent.  535  mm. 

2  per  cent.  1016  mm. 
2.74  per  cent.  1518  mm. 
4  per  cent.  2052  mm. 
6  per  cent.  3°75  mm. 
i  per  cent.  535  mm. 


-Pfeffer,— Ames  Sci.  Mem's.) 

Effect  of  Temperature. 
(i%  sol.  of  cane  sugar.) 
Temperature.          Pressure. 
14.2°  C.  510  mm. 

32.0°  C.  544  mm. 


6.8°  C. 
I3-7°  C. 
22.0°  C. 


505  mm. 
525  mm. 
548  mm. 


SOLUTIONS. 


It  will  be  seen  that  Pfeffer  found  the  osmotic  pressure  of  a  i  per  cent, 
sugar  solution  at  6.8°  C.  to  be  equal  to  50.5  cm.  of  mercury  or  50.5  x 
13.59  grm.  per  sq.  cm.  A  i  per  cent,  sugar  solution  contains  approx- 
imately i  grm.  of  sugar  per  100  cm.s  of  solution.  As  the  molecular 
weight  of  sugar  (C12H22OU)  is  342,  then  a  gram-molecular  weight  is 
contained  in  34,200  cm.3  of  the  solution.  This  volume  represents  then 
the  molar  volume  at  6.8°  C.  or  279.8°  C.  on  absolute  scale  (T). 


PV  =  RT, 


PV 


or,  -  =  R  =  84800  gr.  cms. 

T 

50.5  x  13.59  x  34200 

Substituting,—  -  =  83,900  (approx.). 

279.8 

OSMOTIC    PRESSURE  AND    MOLECULAR    WEIGHTS. 

(Sugar  Solutions  at  about  20°  C.) 
(Morse  and  Frazer.) 


Weight 

Volume 

normal 

normal 

Moles  in  1000 

Moles  per 

grm.  H2O. 

Liter. 

(W) 

0.05 

0.04948 

o.  10 

0.09794 

o.  20 

o.  19192 

0.25 

0.23748 

0.30 

0.28213 

0.40 

0.36886 

0.50 

0.45228 

0.60 

0.53252 

0.70 

0.60981 

0.80 

0.68428 

0.89101 

0.75000 

0.90 

0.75610 

1  .00 

0.82534 

Pressures  at  Same 

Temperature.  W(22.4 

M=- 
Gaseous.    Osmotic  (P). 


o.o824t). 


i .  21 

2.40 

4.82 

6.06 

7.22 

9.68 

12.07 

14-58 

17. 16 

19.17 

21.48 

21.73 

24.27 


1.26 
2.44 
4-78 
6-05 

7-23 
9.66 
12.09 
14-38 
17.03 
19-38 

21  .  21 
21. 8l 
24.49 


Mean,    341.2 


Van't  Hoff  summed  up  these  results  in  the  form  of  a  law  which  bears 
his  name:  —  "  The  osmotic  pressure  of  a  substance  in  solution  is  the  same 
pressure  which  that  substance  would  exert  were  it  in  gaseous  form  at  the 
same  temperature  and  occupying  the  same  volume." 

In  a  paper  which  appeared  in  the  Amer.  Chem.  Jour.,  in  July,  1905, 
Morse  and  Frazer  show  as  the  result  of  a  series  of  accurate  experiments, 
that  Van't  Hoff's  law  holds  for  solutions  of  sugar  if  the  words,  "volume 


88  EXPERIMENTAL  CHEMISTRY. 

of  the  pure  solvent,"  are  substituted  for  the  word  "volume,"  which  refers 
to  the  total  volume  of  the  solution.  They  say:  "When  we  dissolved 
a  gram-molecular  weight  of  cane  sugar  (342.22  grm.)  in  1000  grams  of 
water,  i.e.,  in  that  mass  of  solvent  which  has  the  unit  volume,  i  liter,  at 
the  temperature  of  maximum  density,  we  found  its  osmotic  pressure, 
at  about  20°,  in  quite  close  accord  with  the  pressure  which  a  gram- 
molecular  weight  of  hydrogen  would  exert,  at  the  same  temperature, 
if  its  volume  were  reduced  to  i  liter,  i.e.,  to  that  volume  which  the  unit 
mass  of  solvent  has  at  the  temperature  of  greatest  density."  Or  in 
other  words,  a  substance  in  solution  "  exerts  an  osmotic  pressure  through- 
out the  larger  volume  of  the  solution  equal  to  that  which  as  a  gas  it  would 
exert  if  confined  to  the  smaller  volume  of  the  pure  solvent." 

" It  should  be  borne  in  mind,"  says  Walker,  "that  the  osmotic  pressure 
in  a  solution  may  be  regarded  as  always  present,  whether  a  semiper- 
meable  membrane  renders  it  visible  or  not.  The  osmotic  pressure  in 
the  ordinary  reagent  bottles  of  the  laboratory  is  of  the  dimensions  of 
50  atmospheres.  This  pressure  is,  of  course,  not  borne  by  the  walls 
of  the  bottle  nor  is  it  apparent  at  the  free  surface  of  the  liquid.  Where 
the  liquid  comes  in  contact  with  the  enclosing  vessel  there  we  find  a 
liquid  surface,  and  a  consideration  of  the  magnitude  of  the  forces  at  work 
in  the  phenomena  of  surface  tension  leads  us  to  believe  that  the  pressure 
at  right  angles  to  the  free  surface  of  a  liquid,  and  directed  towards  the 
interior  of  the  liquid,  is  measurable  in  hundreds  and  even  thousands  of 
atmospheres.  Osmotic  pressures,  then,  large  as  they  are  in  ordinary 
solutions,  are  small  compared  to  the  surface  pressures  in  liquids,  and 
their  existence  is  consequently  not  evident  at  the  free  surface  of  liquids. 
It  is  only  when  these  surface  pressures  are  got  rid  of  that  we  can  measure 
osmotic  pressures  directly.  The  liquid  solvent  can  easily  penetrate  the 
semipermeable  membrane,  so  at  the  semipermeable  membrane  there  is 
no  surface  pressure  of  the  ordinary  type.  This  continuity  of  the  liquid 
through  the  semipermeable  partition  gives  us,  therefore,  the  opportunity 
of  determining  differences  of  internal  pressure  in  the  solution  and  the 
solvent.  Various  hypotheses  have  been  put  forward  to  explain  the  nature 
of  osmotic  pressure,  but  none  of  them  can  be  accounted  satisfactory." 

Experiment  IV. — Effect  of  Temperature  on  Solubility  of  Solids. 

(a)  Temperature  increases  solubility.  Add  sodium  chloride  to  a 
test  tube  half  filled  with  water;  shake  and  continue  to  add  salt  until  no 
more  will  dissolve.  The  solution  is  said  to  be  saturated  at  the  temper- 
ature of  the  solution.  Heat  the  test  tube  and  contents;  add  salt  until 
solution  is  saturated  at  the  higher  temperature;  set  tube  aside  to  cool 
if  there  is  no  undissolved  salt  in  it,  otherwise,  filter  and  observe  filtrate 
on  cooling.  Is  the  supernatant  liquid  saturated  at  the  prevailing  tem- 
perature? Decant  a  portion  of  the  clear  liquid  into  another  test  tube; 
cool  the  contents  to  a  lower  temperature  by  immersing  tube  in  a  mix- 
ture of  ice  and  water,  or  allow  cold  tap  water  to  drip  upon  tube.  Is 
more  salt  deposited?  Is  it  in  equilibrium  with  the  salt  in  solution? 


SOLUTIONS.  8() 

Can  the  equilibrium  be  destroyed  temporarily  by  adding  water?     Try. 
Explain.     Boil  solution  in  test  tube  until  salt  is  deposited.     Explain. 

(b)  Temperature     diminishes     solubility.     Prepare     a     concentrated 
solution  of  calcium  citrate  at  the  temperature  of  the  laboratory;  heat, 
but  do  not  boil,  as  it  is  desired  to  avoid  the  vaporization  of  any  appre- 
ciable amount  of  water.     Results  ?     Explain  . 

(c)  Repeat  (b)  using  calcium  hydroxide. 

Experiment  V. — Supersaturated  Solutions. 

Fill  a  test  tube  of  medium  size  nearly  full  of  crystallized  sodium  sul- 
phate; add  4  or  5  cm.3  of  water  and  heat  gently  until  solution  has  a  tem- 
perature of  about  30°  C.;  shake;  add  salt  until  a  saturated  solution  is 
procured;  pour  solution  into  a  clean,  dry  test  tube  or  small  flask;  cover 
the  vessel;  allow  it  to  cool,  then  introduce  a  small  crystal  of  sodium 
sulphate.  If  solution  has  been  prepared  properly  the  excess  of  salt 
will  crystallize  out  of  the  solution.  Is  the  supersaturated  solution  a  case 
of  stable  or  unstable  equilibrium  when  in  contact  with  the  solid  solute  ? 

Experiment  VI. — Thermal  Phenomena  Accompanying  the  Dissolving 
of  a  Solute  in  a  Pure  Solvent. 

Measure  10  cm.3  of  water  into  each  of  five  test  tubes;  take  the  tem- 
perature of  the  water  in  each  tube;  to  the  water  in  one  of  the  tubes  add 
slowly  5  grm.  (see  sp.  gr.,  do  not  weigh)  of  concentrated  H2SO4,  stirring 
carefully  with  the  thermometer  as  acid  is  added.  When  the  contents 
of  tube  are  homogeneous,  i.e.,  one  phase,  record  the  reading  of  the 
thermometer;  remove  thermometer  and  clean  it. 

Repeat  above  using  separately,  5  grm.  of  solid  ammonium  chloride 
(NH4C1),  dehydrated  CuSO4,  Na^C^ic  H2O,  Na^SO,.  Tabulate  data. 
Conclusions  ? 

Experiment  VII. — Relative  Solubility  of  Solids.     Effect  of  Temperature. 

(a)  Place  i  grm.  of  CuO  in  a  test  tube;  add  6  cm.3  of  water.     Does  it 
dissolve?     Heat  the  mixture  to  boiling.     Effect? 

(b)  To  5  grm.  of  NaCl  add  6  cm.3  of  water.     Does  the  salt  dissolve? 
Heat  mixture  to  boiling  in  an  endeavor  to  dissolve  all  of  the  salt.     Are 
you  successful?     Pour  the  hot  saturated  solution   upon   a  dry  filter; 
collect  the  filtrate  in  a  test  tube  and  cool.     Results?     Explain. 

(c)  Pulverize  about  15  grm.  of  K2Cr2O7;  to  5  grm.  in  a  test  tube  add 
6  cm.3  of  water.     Will  all  of  the  solid  dissolve  in  the  solvent  ?     Heat 
mixture  to  boiling.     Has  all  of  the  solid  dissolved?     Cool  the  solution. 
Observe  the  effect  of  lowering  the  temperature. 

Note.— The  pulverized  K2Cr2O7  will  be  needed  in  Exp.  VIII. 

Experiment.  VIII. — (Quant.)  Prepare  a  saturated  solution  of  K2Cr2O7; 
to  8  grm.  of  the  powdered  substance  in  a  flask  or  beaker,  add  50  cm.3  of 
distilled  water;  assist  the  process  of  dissolving  by  frequently  shaking  it. 
If  the  directions  have  been  adhered  to  the  solution  will  probably  be 


90  EXPERIMENTAL  CHEMISTRY. 

saturated  at  the  end  of  10  min.  Take  the  temperature  of  the  solution. 
Into  a  weighed  evaporating  dish,  weigh  accurately  20-30  grm.  of  the 
solution;  evaporate  to  dryness;  cool;  weigh.  Repeat  heating  and  weigh- 
ing. Calculate  the  weight  of  dichromate  in  i  liter  of  the  solution  (satur- 
ated) at  the  observed  temperature.  How  many  gram-molecules  (moles) 
of  potassium  dichromate  in  a  liter  of  the  saturated  solution  ? 

Terminology  of  Solutions. — When  a  substance  dissolves  in  a  liquid 
there  is  for  each  temperature  and  pressure  a  definite  solubility,  i.e.,  a 
definite  relation  between  the  solute  and  the  solvent.  If  the  solution  con- 
tains less  of  the  solute  than  corresponds  to  the  latter's  solubility  in  the 
solvent,  it  is  unsaturated;  if  the  amount  in  solution  is  in  excess  of  the 
amount  required  to  saturate  it,  the  solution  is  said  to  be  supersaturated. 
The  test  of  the  degree  of  saturation  of  a  solution  is  made  by  placing  a 
portion  of  the  solid  solute  in  contact  with  it;  if  the  solution  is  unsaturated, 
a  portion  of  the  solute  will  dissolve;  if  supersaturated,  a  portion  of  the 
solute  will  separate  from  the  solution,  and  continue  to  be  deposited  until 
the  solution  is  saturated  and  a  condition  of  stable  equilibrium  is  estab- 
lished. 

Although  the  above  nomenclature  is  convenient  it  does  not  convey 
the  definite  information  which  is  so  much  desired  by  the  chemist.  Be- 
cause of  this,  the  concentrations  of  solutions  are  frequently  expressed  in 
terms  of  physical  or  chemical  units. 

The  concentration  of  any  substance  is  the  total  amount  of  that  sub- 
stance in  solution  in  a  unit  volume. 

The  solubility  of  a  substance  is  expressed  in  terms  of  the  number  of 
grams  or  gram-molecules  which  can  be  dissolved  in  a  unit  volume  at  a 
given  temperature  and  pressure. 

A  solutio'n  is  referred  to  as  being  standard  when  its  concentration  is 
known.  Of  more  frequent  use,  however,  are  the  terms,  normal  solutions 
and  molar  solutions. 

A  normal  solution  is  a  standard  solution  which  contains  in  one  liter  the 
hydrogen  equivalent  of  the  active  reagent,  expressed  in  grams.  Thus 
a  normal  solution  of  HC1  contains  36.45  grams  of  hydrogen  chloride; 

1TT0^     Mol.  Wt.                          ,  ™TT    Mol.Wt.  . 

normal  H2SO4,—  -  grams;  normal  KOH, grams;  normal 

iodine,  126.97  grams  in  a  liter.  If  a  liter  contains  TV  of  an  equivalent 
weight  it  is  designated  tenth  or  deci-normal  (.iN);  if  TJ~g-  of  an  equivalent, 
hundredth  or  centi-normal  (.oiN). 

A  molar  solution  is  a  standard  solution  which  contains  one  mole  or  one 
gram-molecular  weight  of  the  solute  in  one  liter  of  solution.  Thus  a 
molar  solution  of  HC1  contains  36.45  grm.  of  hydrogen  chloride;  molar 
H2SO4,  98.07  grm.;  molar  NaOH,  40.05  grm.  of  sodium  hydroxide. 

Influence  oj  Temperature  On  Solubility  oj  Solids. — It  is  a  general  rule, 
if  the  temperature  changes,  the  solubility  changes.  The  solubility  of 
the  majority  of  solid  substances  increases  with  increase  of  temperature; 
however,  there  are  cases  in  which  the  solubility  is  decreased  by  rise  of 


SOLUTIONS. 


temperature.  The  increase  in  the  solubility  of  potassium  chloride  is 
approximately  proportional  to  the  increase  of  temperature.  The  solu- 
bility of  calcium  citrate  is  greater  at  10°  C.  than  it  is  at  higher  temper- 
atures, say  70°  C.;  the  same  is  true  of  calcium  hydroxide.  It  is  quite 
probable  that  in  the  majority  of  cases  there  is  a  fall  of  temperature  due 
to  the  mere  act  of  solution,  but  the  heat  of  the  chemical  combination 
which  undoubtedly  follows  the  process  of  dissolving  in  many  instances, 
is  frequently  so  much  in  excess,  that  heat  alone  is  the  observed  result 
of  solution.  For  example,  when  sulphuric  acid  is  dissolved  in  water, 
it  combines  with  the  water  to  form  hydrates  of  sulphuric  acid  and  a  large 
quantity  of  heat  is  developed;  the  same  conduct  is  displayed  by  dehy- 
drated salts,  like  Na2CO3  and  N^SO^  The  question,  is  there  a  relation 
between  the  influence  of  temperature  upon  solubility,  and  some  other 
property  of  substances,  may  be  answered  by  an  application  of  that 
principle  which  is  the  basis  of  explanation  of  all  influences  affecting 
equilibrium;  the  principle  referred  to  is  the  one  enunciated  in  the  form 
of  La  Chatelier's  Theorem.  If  a  substance  dissolves  with  an  absorp- 
tion of  heat,  its  solubility  will  increase  with  rise  of  temperature;  if,  on  the 
other  hand,  heat  is  developed  on  solution,  solubility  will  decrease  with 
lise  of  temperature.  When  substances  dissolve  without  thermal  alter- 
ations, solubility  is  practically  independent  of  temperature.  The 
solubility  of  sodium  chloride  is  affected  but  little  by  temperature  alter- 
ations. A  simple  test  of  the  influence  of  temperature  upon  the  solubility 
of  a  given  substance  in  a  particular  solvent  may  be  made  by  preparing 
a  saturated  solution  at  a  given  temperature,  then  observe  if  increase  in 
temperature  causes  deposition  of  the  solid;  if  solute  is  not  deposited, 
add  a  small  crystal  of  it  to  the  solution;  in  case  of  supersaturation  the 
solid  will  deposit  from  the  solution,  and  if  the  solution  is  now  unsaturated, 
as  the  result  of  the  rise  in  temperature,  the  crystal  will  dissolve. 


TABLE  OF  SOLUBILITY  OF 
(Grams  dissolved  by  100  cm.3 


SOLIDS. 
of  water.) 


At  o° 

Sodium   Chloride 35.6 

Sodium  Nitrate 72.9 

Sodium     Sulphate      (hy- 

drated) 5.02 

Potassium  Chloride 30.0 

Potassium   Nitrate 13.3 

Potassium  Sulphate 8.46 

Potassium  Bichromate  .  .  4.9 

Ammonium  Chloride.  .  .  .  28.4 

Copper  Sulphate 18.2 

Calcium  Hydroxide 0.174 

Calcium  Sulphate 0.205 

Calcium  Chloride 49-59 


C. 

grm. 

« 


35-63 
87-5 


grm. 


100°  C. 
39.9      grm. 
180.0 

(Parts  dissolved 
by  100  parts  of 
80%  alcohol 
at  15°.) 

1.22 

2.8 

42.4 


(32.38°  C.)       (anhydrous) 


i. 3  (sp.gr.  .94) 


34.7 

56.6 

29.0 

247.0 

(18° 

-•) 

10.9 

26.2 

1  3.  i 

102.0 

37.28 

72.80 

42.31 

203.22 

0.13 

0.08 

0.208 

.218 

74.00 

'     .    I49-98 

0.45  (70%) 
0.4 

0.21    (40%) 

12.0    (abs.8°) 


60.00  (abs.  80°) 


92  EXPERIMENTAL  CHEMISTRY. 

Heat  of  Solution. — We  have  observed  that  the  process  of  solution  is 
usually  accompanied  by  thermal  phenomena,  i.e.,  heat  is  absorbed  or 
evolved.  This  is  called  the  heat  of  solution.  It  is  obvious  that  for  differ- 
ent amounts  of  water,  this  will  vary;  in  order  to  establish  uniformity, 
the  heat  of  solution  is  now  usually  understood  to  be  the  amount  of  heat 
liberated  or  absorbed  by  the  solution  of  i  gram-molecule  of  the  sub- 
stance under  consideration,  in  such  a  large  quantity  of  water  (solvent) 
that  the  further  addition  of  water  will  not  yield  an  additional  heat  effect. 
The  addition  of  the  first  quantity  of  water  to  the  solute  produces  a  rela- 
tively larger  thermal  change  than  the  succeeding  additions  of  equal 
amounts  of  the  solvent,  therefore  the  use  of  an  unlimited  volume  of  the 
solvent  as  suggested  above. 

Ostwald  defines  the  "heat  of  solution"  as  "the  heat  which  is  taken 
up  or  given  out  when  a  further  quantity  of  salt  is  dissolved  in  a  solution 
saturated  at  a  definite  temperature."  "This  quantity  of  heat,"  he  says, 
"must  not  be  confused  with  that  which  accompanies  the  solution  of  a 
salt  in  the  pure  solvent,  and  which  is  usually  what  is  measured.  In 
the  case  of  difficulty  soluble  substances,  it  is  true,  the  two  are  not  greatly 
different;  but  where  the  substances  are  soluble  in  large  amounts,  they 
can  have  not  only  a  different  value  but  even  a  different  sign."  Richards, 
T.  W.,  defines  the  "heat  of  solution,"  of  a  substance,  as  that  which  is 
developed  or  absorbed  when  the  final  quantity  (i  mole)  of  salt  which 
finishes  the  saturation  of  an  unlimited  quantity  of  water,  is  added.  An 
erroneous  application  of  this  principle  has  led  to  many  apparent  and 
perplexing  contradictions. 

The  following  tables  give  the  heat  of  solution  of  a  number  of  substances 
in  a  large  quantity  of  water  at  i8°-2o°  C.  This  is  designated  by  the 
addition  of  the  abbreviation  Aq  (aqua)  to  the  formula  or  symbol  of  the 
substance: 

HEAT    OF    SOLUTION    OF    SOLIDS. 
(From  Hortsmann,  Theoret.  Chem.,  p.  502.) 
NaOH,  Aq  =  +9780  cal.  KOH,  Aq  =  +12500  cal. 

NaCl,  Aq  =  — 1180  cal.  KOH.2H2O,  Aq  -  —30  cal. 

)3,  Aq  =  +5640  KC1,  Aq  =  — 4440 

EO  H2O,  Aq  =  —16160  K2SO4,  Aq  =  —6380 

)4",  Aq  =  +460  NH4C1,  Aq  =  —3880 

o  H2O,  Aq  =  —18760  LiCl,  Aq  =  +8440 

NaC2H3O2,  Aq  -  +4200  CaCl2,  Aq  =  +3258 

Na  Br,  Aq  =  — 190  AgCl,  Aq  =  — 15800 

Na  Br^HjO,  Aq  =  — 4710  Ag  Br,  Aq  =  — 20200 

C12H22On,  Aq  =  —800  Ag  I,  Aq  =     —26600. 

HEAT   OF    SOLUTION    OF    LIQUIDS. 

H2SO4,       Aq  =  +17800  cal.  CH3OH,    Aq  =  +2000  cal. 

C2H4O2,     Aq  =  +420  cal.  C2H5OH,  Aq  =:  +2540  cal. 

(C2H5)2O,  Aq  =  +5940  cal.  C3H7OH,  Aq  =  +3050  cal. 


SOLUTIONS.  93 

HEAT    OF    SOLUTION    OF    GASES. 

H9F9,  Aq  =  +11800  cal.*  Cl.,,     Aq  =••  +4870  cal. 

HC1,   Aq  =:  +17310  cal.  CO2,  Aq  =  ••  +5880  cal. 

HBr,  Aq  =••  +19940  cal.  NH3,  Aq  =  +8430  cal. 
H  I,    Aq  =  +19210  cal. 

All  gases  which  have  been  examined,  and  all  liquids,  as  a  rule,  dis- 
solve with  evolution  of  heat,  but  solid  substances  show  no  such  uniformity 
of  conduct.  Some  of  them,  solids,  dissolve  with  an  absorption  of  heat 
while  others  behave  in  a  contrary  manner.  Nernst  suggests  that  the 
explanation  of  this  is  simple.  Assume  that  a  given  substance  in  the 
gaseous  condition  always  dissolves  with  an  evolution  of  heat,  then  if  it  is 
reduced  to  the  liquid  state  before  it  is  placed  in  solution,  it  will  dissolve 
with  an  evolution  or  absorption  of  heat  according  as  its  heat  of  vapori- 
zation is  less  or  greater  than  its  heat  of  solution.  The  sign  of  the  heat 
of  solution  of  the  substance  in  the  solid  phase  will  depend  upon  the 
relative  magnitude  of  its  heat  of  sublimation  and  its  heat  of  solution 
when  in  the  gaseous  phase.  Interpreting  the  data  presented  by  the  above 
tables  in  the  light  of  the  above  suggested  explanation,  we  conclude  that 
the  "  heat  of  solution  in  the  gaseous  state  is  always  greater  than  the  heat 
of  vaporization,  but  it  is  usually  smaller  than  the  heat  of  vaporization 
plus  the  heat  of  fusion,  i.e.,  it  is  smaller  than  the  heat  of  sublimation." 

Effect  of  Pressure  on  Solubility  oj  Solids. — A  solution  usually  occupies 
less  space  than  the  sum  of  its  components;  it  never  occupies  more  than 
the  sum.  In  some  cases,  for  example,  solutions  of  Li(OH),  Ba(OH)2, 
NiSO4,  CaSO4,  the  volume  of  the  solution  is  less  than  the  original  volume 
of  the  solvent.  Inasmuch  as  the  process  of  solution  is  frequently  accom- 
panied by  small  changes  in  volume,  it  is  evident  that  if  a  solution  is  in 
equilibrium  with  the  solid  substance,  pressure  will  alter  the  solubility 
in  a  manner  which  may  be  anticipated  by  the  application  of  the  principle 
enunciated  in  Le  Chatelier's  Theorem.  The  effect  of  pressure  is  so 
small  that  pressures  equivalent  to  hundreds  of  atmospheres  are  required 
to  effect  a  change  in  solubility  so  small  that  it  can  scarcely  be  measured. 

The  question  may  arise,  what  is  the  explanation  of  these  volume- 
changes  observed  while  preparing  solutions;  why  does  alcohol  (ethyl) 
and  water  when  mixed  in  equal  proportions  by  volume,  contract  about 
3  per  cent,  of  the  total  volume?  Why  do  10,000  cm.3  of  water  and  27.5 
cm. 3  (58.5  grm.)  of  sodium  chloride  yield  a  solution  whose  volume  is 
about  10016.5  cm. 3  ?  The  question  can  not  as  yet  be  answered  definitely. 
It  has  been  suggested  that  the  "associated"  molecules  of  the  water  are 
broken  up  during  the  dissolving  of  the  solute,  and  as  the  molecules 
occupy  less  space  when  in  a  "  non- associated "  condition,  this  may  ac- 
count for  the  observed  contraction.  Again,  it  may  be  due  to  the  fact 
that  the  dissolved  substance  undergoes  hydration,  i.e.,  attracts  water 
to  itself,  forming  a  hydrate.  Now  the  water  in  hydrated  salts 
occupies  less  space  than  when  not  thus  combined.  This  may  partially 
*  580  cal.  (ext.  work)  should  be  subtracted  from  above  figures. 


94  EXPERIMENTAL  CHEMISTRY. 

explain  the  conduct  of  the  solutions  referred  to  above  as  examples  in 
which  the  volume  of  the  solution  was  less  than  original  volume  of  solvent. 
Just  where  the  contraction  is  no  one  seems  to  know. 

Solubility  of  Liquids. — It  is  frequently  convenient,  when  speaking  of 
the  relative  solubility  of  liquids  in  liquids,  to  refer  to  the  solubility  of 
"  pairs  of  liquids. "  These  groups  of  two  many  be  divided  into  four  orders : 

SOLUBILITY    OF    PAIRS    OF    LIQUIDS. 

i. — Miscible  in  all  proportions. 

2. — Partially  miscible. 

3. — Immiscible  (mutually  insoluble). 

4. — Pairs  which  are  in  one  order  at  some  temperatures  but  in  another 
order  at  other  temperatures. 

The  first  order  embraces  those  pairs  in  which  the  solubility  of  the 
two  liquids  in  each  other  is  unlimited;  the  second  order,  those  pairs  in 
which  the  solubility  of  the  two  liquids  in  each  other  is  limited;  the  third 
order,  those  pairs  in  which  the  two  liquids  are  not  wholly  but  practically 
insoluble  in  each  other,  as  mercury  and  wrater,  organic  liquids  and  water, 
etc.;  the  fourth  order,  explains  itself,  and  suggests  that  orders,  one,  two 
and  three,  will  pass  gradually  from  one  to  the  other  with  sufficient  change 
of  temperature. 

Experiment  IX. — Relative  Solubility  of  Liquids. 

(a)  To  5  cm.3  of  water  in  a  test  tube  add  a  few  drops  of  alcohol;  shake. 
How  many  phases  are  present?     Do  alcohol  and  water  mix?     Add 
alcohol  in  small  quantities  until  5  cm.3  of  alcohol  have  been  added? 
Does  the  mixture  become  homogeneous  on  shaking  after  each  addition 
of  alcohol?     To  which  order  does  the  above  pair  of  liquids  belong, 
basing  your  answer  on  observed  results  ? 

(b)  Repeat  (a)  using  5  cm.3  of  ether. 

(c)  Repeat  (a)  using  2  cm.3  of  kerosene. 

(d)  Repeat  (a),  using  2  cm.3  of  benzol  or  toluol.     Define  "phase." 

Experiment  X. — Mutual  Solubility  of  Liquids. 

To  3  cm.3  of  distilled  H2O  in  a  test  tube  add  3  cm.3  of  ether  and  shake 
vigorously.  How  many  phases  are  present  ?  Does  each  layer  represent 
a  solution?  Set  tube  aside — to  be  used  later.  To  3  cm.3  of  ether  add 
3  cm.3  of  distilled  H2O.  Proceed  as  above.  Compare  the  contents  of 
the  two  test  tubes.  Pour  the  contents  of  one  tube  into  the  other  tube, 
thereby  mixing  the  solutions;  allow  tube  to  stand  until  there  are  two 
distinct  phases,  i.e.,  two  layers.  Dehydrate  about  a  grm.  of  pulverized 
CuSO4.io  H2O,  the  salt  should  be  nearly  white.  Now  devise  a  method 
for  proving  that  the  top  layer  of  the  solution  in  the  test  tube  is  a  solution 
of  water  in  ether.  Also  prove  that  ether  is  present  in  the  lower  layer 
which  is  a  solution  of  ether  in  water. 

Hint. — When  a  solution  of  ether  and  water  is  gently  heated  in  a  test 
tube,  the  ether  is  rapidly  expelled  and  can  be  inflamed. 


SOLUTIONS.  95 

The  heavier  liquid  at  the  bottom  of  the  test  tube  is  an  aqueous  solution 
of  ether  containing  about  10  per  cent,  of  ether;  the  upper  layer  is  an 
etheral  solution  of  water  containing  about  3  per  cent,  of  water.  A 
similar  phenomenon  is  exhibited  by  the  use  of  a  concentrated  solution 
of  potassium  carbonate  and  aqueous  ammonia,  sp.  gr.  0.88.  This  is 
probably  the  only  known  case  in  which  aqueous  solutions  of  inorganic 
substances  behave  in  this  manner. 

Experiment  XI. — Effect  of  Temperature.  Critical  Solution  Tempera- 
ture. 

(a)  Increase  in  temperature  produces  complete  miscibility.     To  a  test 
tube  containing  about  3  cm.3  of  solid  or  liquid  phenol*  add  about  10 
cm.  s  of  water;  cork  the  tube  and  shake  vigorously;  a  milky  appearing 
mixture,  an   emulsion,  is   the  result;  allow  the  tube  to  stand  for  some 
time,  or  better,  remove  the  cork  and  heat  gently — when  the  mixture  will 
separate  into  layers — an  aqueous  solution  of  phenol  above,  and  a  solu- 
tion of  water  in  phenol  below;  now  place  the  test  tube  in   the    flame 
and  heat  gradually  until  the  line  of  demarcation  between  the  two  phases 
slowly  disappears;  warm  a  thermometer  until  its  reading  is  about  80°  C., 
then  introduce  it  into  the  test  tube  so  that  its  bulb  will  be  immersed 
in  the  contents,  and  determine  the  temperature  at  which  the  mixture 
becomes  homogeneous;  cool  the  tube  and  note  the  reading  of  the  ther- 
mometer when  the  milky  appearance  is  observed.     Take  the  average 
of  the  two  thermometer  readings.     What  are  your  conclusions  as  to  the 
mutual  solubility  of  this  pair  of  liquids  at  the  temperature  at  which  the 
contents  became  homogeneous?     At  the  temperature  of  the  laboratory 
(20°  C.)? 

(b)  Lowering  of  the  temperature  produces  complete  miscibility.     Use 
dimethylamine  and  water.     (Instructions.) 

The  lowest  temperature  at  which  pairs  of  liquids  like  phenol  and 
water  become  miscible  in  all  proportions  is  called  the  "critical  solution 
temperature." 

Experiment  XII. — Distribution  of  Solute  Between  Two  Immiscible 
Solvents.  Law  of  Distribution. 

To  a  small  flake  of  sublimed  iodine  in  a  test  tube  add  10  cm.3  of 
distilled  H2O;  shake  it  vigorously  for  several  minutes  as  the  iodine  is 
nearly  insoluble  and  dissolves  slowly;  pour  off  the  clear  solution  into  a 
clean  test  tube.  Repeat  above  operation  using  an  aqueous  solution  of 
KI  as  the  solvent.  Divide  each  solution  into  two  portions;  to  one  portion 
of  each  solution  add  3  cm.3  of  ether;  and  to  the  other  portion  of  each, 
add  3  cm.3  of  carbon  disulphide;  shake  the  tubes.  Results?  In  each 
case,  after  adding  ether  or  carbon  disulphide,  how  many  phases  were 
present  in  the  respective  test  tubes?  Drawing  your  conclusions  from 

*  Caution: — A  very  small  quantity  of  phenol  (carbolic  acid,  C6H6O)  will  attack  the 
flesh.  Handle  vessel  containing  it  with  the  utmost  care. 


96  EXPERIMENTAL  CHEMISTRY. 

observed  results,  arrange  the  solvents  in  the  order  in  which  the  increasing 
solubility  of  iodine  is  shown. 

Law  of  Distribution. — The  solute  is  distributed  between  the  two 
solvents  in  such  a  way  that  the  ratio  oj  its  concentration  in  each  is  a  con- 
stant. This  constant  is  practically  independent  of  the  absolute  con- 
centration and  is  dependent  only  upon  temperature  and  the  nature  of 
the  solute  and  the  two  solvents.  (It  is  obvious  that  the  ratio  of  the 
concentrations  is  the  ratio  of  the  solubilities  in  the  separate  solvents.) 
Also,  if  two  or  more  substances  are  placed  in  solution  the  coefficients  of 
distribution  are  the  same  as  if  each  substance  were  present  alone.  (This 
reminds  one  of  the  analogous  conduct  of  gases.)  In  connection  with 
the  above  experiment,  it  may  be  noted  that  in  i  cm.3  of  the  carbon 
disulphide  solution  there  will  be  found  600  times  as  much  iodine  as  in 
i  cm.3  of  the  aqueous  solution  of  iodine. 

Experiment  XIII.— Pressure  of  Gas  Affects  Solubility. 

Half  fill  the  large  test  tube  which  you  used  in  generating  oxygen,  with 
water;  cork  the  tube  with  a  rubber  stopper  perforated  with  a  single  hole; 
shake  the  tube  until  the  water  is  saturated  with  air;  connect  tube  with 
an  air  pump  and  exhaust  the  air  from  above  the  solution;  observe  the 
small  bubbles  of  gas  passing  out  of  the  liquid.  Explain.  What  would 
be  the  probable  effect  of  increasing  the  pressure  of  the  gas  ? 

Henry's  Law. — At  a  constant  temperature,  a  given  quantity  of  the  liquid 
solvent  will  dissolve  weights  of  a  gas  which  are  proportional  to  the 
pressure  of  the  gas.  This  law  holds  when  the  gases  are  only  moder- 
ately soluble.  When  a  mixture  of  gases  dissolves  in  a  liquid,  each  dis- 
solves as  if  it  were  present  alone — i.e.,  each  dissolves  according  to  its 
own  partial  pressure.  In  other  words,  the  gas  distributes  itself  between 
the  liquid  and  gaseous  phases  according  to  a  constant  (K)  which  is  a 
characteristic  of  each  gas.  (Recall  the  Law  of  Distribution.) 

Experiment  XIV. — Effect  of  Solvent  on  Solubility. 

Half  fill  a  test  tube  with  water;  observe  that  it  is  perfectly  clear;  add 
5  cm.3  of  alcohol.  Explain  the  liberation  of  the  small  bubbles  of  air. 

SOLUBILITY    OF    GASES. 
(o°  C.  760.  mm.) 

Name  of  gas.  i  Vol.  of  H2O       i  Vol.  of  C2H0O 

Dissolves,  Dissolves, 

Hydrogen 0.02  vols.  0.07  vols. 

Oxygen 0.04  vols.  6.28  vols. 

Nitrogen 0.02  vols.  0.13  vols. 

Carbon  dioxide 1.79  vols.  4.32  vols. 

Hydrogen  sulphide 4.37  vols.  17.9    vols. 

Ethylene 0.25  vols.  3.6    vols. 

Ammonia 1148.80  vols. 

Hydrochloric  acid 505,00  vols. 


SOLUTIONS.  97 

Experiment  XV. — Effect  of  Temperature  on  Solubility. 

Half  fill  a  1000  cm.3  flask  with  water;  cork  it  and  shake  it  until  water 
is  saturated  with  air;  determine  its  temperature.  Calibrate  a  300  cm.3 
flask;  provide  it  with  a  rubber  cork  through  which  passes  a  delivery  tube 
which  will  dip  beneath  the  surface  of  the  water  in  the  pneumatic  trough; 
fill  the  300  cm.3  flask  and  delivery  tube  with  the  water  saturated  with 
air;  heat  the  flask  to  boiling;  collect  the  air  which  is  liberated  in  an  in- 
verted eudiometer  by  water-displacement;  read  the  barometer,  determine 
the  temperature  of  the  water  in  the  trough,  equalize  the  water  levels 
and  take  the  reading  of  the  eudiometer.  Calculate  the  volume  of  gas 
dissolved  by  a  cm.3  of  water  under  the  conditions  which  prevailed  while 
preparing  the  aqueous  solution  of  air. 

Note. — A  test  tube  may  be  used  in  place  of  the  eudiometer  by  marking 
the  volume  of  gas  with  gummed  paper,  and  then  calibrating  the  tube. 

SOLUBILITY    OF    GASES. 
(Vols.  of  gas  which  i  vol.  of  H2O  will  dissolve.) 


L/J.O.      \Ji      gClO      V 

Hydrogen. 

V  JJUt~J.Jl      4.       VW1.      VJ 

Oxygen. 

'i    j-j-ovy    win    vj.1* 

Nitrogen. 

Carbon  dioxide 

.0215  vol. 

.0411  vol. 

.0203  vol. 

.0179  vol. 

.0206  vol. 

.0362  vol. 

.0179  vol. 

.0144  vol. 

.0198  vol. 

.0325  vol. 

.0160  vol. 

.0118  vol. 

.0185  vol. 

.02.83  vol. 

.0140  vol. 

.0090  vol. 

5° 
10° 

20° 

The  quantity  of  gas  which  a  liquid  is  capable  of  dissolving  depends 
upon,  (i)  nature  of  solvent;  (2)  nature  of  gas;  (3)  the  temperature  of 
the  solvent;  (4)  the  partial  pressure  of  the  gas. 

Gases  with  reference  to  each  other,  are  regarded  as  being  miscible 
in  all  proportions.  Inasmuch  as  solids  and  liquids  directly  vaporize 
into  gases,  the  latter  are  considered  as  solvents  of  the  former. 

Experiment  XVI.— Lowering  the  Vapor  Tension  or  Elevation  of  the 
Boiling  Point. 

Determine  the  boiling  point  of  distilled  water — proceed  as  in  a  former 
experiment.  In  this  case  determine  the  temperature  of  the  liquid  as 
well  as  that  of  the  vapor.  Record  your  data.  Add  a  sufficient  quantity 
of  CaCl2  to  nearly  saturate  20-25  cm.3  of  the  distilled  water;  heat  the 
solution  to  boiling.  Observe,  (a)  the  temperature  of  the  solution;  (b) 
the  temperature  of  the  vapor.  Record  data.  Explain.  Will  vegetables 
cook  more  rapidly  in  a  boiling  salt  solution  than  in  boiling  distilled  water  ? 
Why? 

Caution. — Suspend  the  thermometer  when  taking  temperatures;  do 
not  allow  it  to  touch  the  flask. 

Experiment  XVII. — Depression  of  the  Freezing  Point. 

Half  fill  a  test  tube  with  distilled  water;  determine  its  freezing  point; 
add  NaCl  until  the  water  is  nearly  saturated;  determine  the  freezing 

7 


98 


EXPERIMENTAL  CHEMISTRY. 


point  of  the  solution.  Has  the  freezing  point  of  the  water  been  de- 
pressed? How  much  (express  in  degrees0  C.)?  (See  Fig.  21,  Beckmann 
apparatus.) 

Elevation  of  the  Boiling  Point.  —  It  has  been  known  for  many  years  that 
the  normal  boiling  or  freezing  point  of  a  liquid  may  be  altered  by  dissolv- 

ing a  quantity  of  a  solid  substance  in  it.  We 
have  observed  that  pure  water  freezes  at  o°  C. 
and  boils  at  100°  C.  (760  mm.),  whereas  the 
freezing  point  of  solutions  of  solids  is  below, 
and  the  boiling  point  above,  normal.  If  the 
boiling  point  of  a  solvent  is  elevated,  it  is 
obvious  that  its  vapor  tension  is  lowered  by  the 
presence  of  the  solute. 

It  has  been  shown  experimentally  that  a 
definite  relation  holds  between  the  freezing  and 
boiling  points  of  a  liquid  and  the  quantity  and 
nature  of  the  dissolved  substance.  Further, 
the  degree  of  the  depression  of  the  freezing 
point  or  the  elevation  of  the  boiling  point  de- 
pends upon  four  factors:  (i)  the  nature  of  the 
solute;  (2)  the  quantity  of  the  solute;  (3)  the 
nature  of  the  solvent;  (4)  the  quantity  of  the 
solvent.  To  make  an  application  —  i  gram  of 
salt  dissolved  in  100  grams  of  acetic  acid  pro- 
duces a  depression  of  the  freezing  point  which 
is  different  from  that  produced  by  dissolving 
the  same  quantity  of  salt  in  100  grams  of  water. 
One  mole  of  any  substance  dissolved  in  100 
grams  of  solvent  must  always  produce  a  cer- 
tain definite  increase  in  the  boiling  point  of 
that  solvent,  because  it  produces  a  definite 
depression  of  the  vapor  tension.  (Recall  rela- 
tion of  osmotic  pressure  to  vapor  pressure.) 

The  elevation  produced  by  i  gram  of  sub- 
stance in  100  grams  of  solvent  is  spoken  of  as 
the  specific  elevation  (d).  It  has  been  found 
that  the  specific  elevation  (d)  produced  by  any 
substance  in  any  given  solvent  when  multiplied 
by  the  molecular  weight  (M)  of  the  solute,  gives 
a  product  which  is  practically  a  constant  (K) 
for  that  particular  solvent.  K  represents  the 

molecular-increase  of  the  boiling  point,  i.e.,  "  that  due  to  the  solution  of 
i  mole  of  substance  in  100  grams  of  solvent,  which  must  be  constant 
for  all  substances  in  the  same  solvent." 


FIG.  21. 


K  =  dM, 


„-£ 


SOLUTIONS.  99 

One  mole  of  any  substance  dissolved  in  100  grams  of  ether  increases 
the  boiling  point  of  the  latter  21.10°  C. 

Ostwald  (Physico-Chemical  Measurements)  gives  the  following 
values  for  the  constant  K  which  is  equal  to  100  K.  :  — 

Ethyl  ether  ..............    2110  Ethyl  acetate  .........  2610 

Benzene  ................    2670  Acetone    .............  1670 

Chloroform  ..............   3660  Water     ..............  520 

Carbon  disulphide  ........    2370  Ethylene  dibromide  .  .  .  6320 

Acetic  acid  ..............    2530  Aniline     .............  3220 

Ethyl  alcohol  ............    1150  Phenol    ..............  3040 

The  following  expression  has  been  developed  for  the  purpose  of 
ascertaining  molecular  weights: 


lOOg 

g 
M  -  100  K  ---  , 

AG 

where  G  represents  the  number  of  grams  of  solvent;  g,  the  grams  of 
solute;  A  the  elevation  of  the  boiling  point  in  degrees,  and  M,  the  molec- 
ular weight  of  the  solute.  The  value  of  K  may  also  be  found  by  a  proc- 
ess of  reasoning  based  upon  thermo-dynamical  principles. 

Beckman  while  determining  the  molecular  weight  of  iodine  in  ether 
recorded  the  following  data: 

K  =  21.10,     g  =  2.0579,     G  =  30.14,      A=  0.566; 
M  =  100  K 

substituting  in  above  equations,— 


M  =  100  K . 

AG 


2.0579 

M   =    IOO  X   2I.IO 


0.566  x  30.14 

M  =  254, 
I2   =  254- 

Morgan  suggests  the  use  of  the  following  simple  proportion  in  cal- 
culating the  molecular  weight: 

2.0579 
21. i  :  0.566  ::  M  grm.  per  100  grm.  :  —        —  x  100. 

30.14 
M  =  254. 


IOO  EXPERIMENTAL  CHEMISTRY. 

Depression  oj  the  Freezing  Point.  —  Using  a  similar  method  of  cal- 
culating experimental  data,  we  are  able  to  ascertain  the  molecular 
weight  of  the  solute  by  observing  the  depression  of  the  freezing  point  of 
the  solvent.  In  1887  Raoult  found  that  "  i  mole  oj  any  substance  dis- 
solved in  100  grams  oj  any  one  solvent  causes  a  constant  depression  oj 
the  freezing  point."  It  will  be  observed  that  the  relation  of  the  solute 
to  the  freezing  point  is  very  similar  to  its  relation  to  the  boiling  point. 
The  mathematical  form  of  the  expression  used  to  calculate  the  molecular 
weight  is  identical  with  the  above.  If  d  equals  the  specific  depression 
and  K  the  molecular  depression,  then,  — 

K 

K  =  dM,        M  =  —  . 
d 

AG 


lOOg 

g 

M  =  100  K  --  , 

AG 

where  G  represents  the  number  of  grams  of  solvent;  g,  the  grams  of 
solute;  A  the  depression  of  the  freezing  point  in  degrees,  and  M  the 
molecular  weight  of  the  solids. 

SOLVENTS. 

Acetic  Acid.  Water. 

Solute.          d.         M.  K.  Solute.      d.          M.  K. 

SO2          0.6015      64  38.5  NH3       1-1705      17-  18.80 

CS2          0.5050     76  38.4  CuSo4    0.1153    159.7  18.41 

CH  C13    0.3247    119.3  38.8  C2H4O2o.3i62      60,  18.97 


Aver.  38.6  Aver.  18.72 

(Ostwald),  38.8  (Ostwald),  18.90 

(Jones),  18.50 

SOLUTIONS    IN    WATER. 

(Jones,  Phys.  Chem.) 

Solute.                   Molecular.                   Solute.  Molecular. 

(Organic.)  Lowering  (K).             (Inorganic.)  Lowering  (K). 

Methy  alcohol                17.3  Hydrochloric  acid  39.1 

Malic  acid                      18.7  Sodium  oxalate  13.2 

Acetone                            17.1  Sodium  hydroxide  36.2 

Acetic  acid                     19.0  Sulphuric  acid  38.2 


SOLUTIONS.  101 

RAOULT'S  LAW  FOR  SOLVENTS. 
(Raoult's  Tables). 

M  K  K/M 

Solvents.  Mol.  Weight.    Mol.  Lowering.    Lowering  produced  by 

i  molecule  in  100  mole- 

cules. 

Formic  acid  46  29  o°-63 

Acetic  acid  60  39  °°-65 

Benzene  78  ,5°  o°.64 

Ethylene  bromide      123  *v'-T5  '  ?          •   '^       °°-59 

Water  18          ,  47  2°.6i 


Raoult's  Law.  —  "//  one  mohci'le  '  of  *  "any  ~ytel)s&&(?e  $s  dissolved  in 
one  hundred  molecules  of  any  liquid  of  a  different  nature,  the  lowering  of 
the  freezing  point  of  this  liquid  is  always  nearly  the  same,  and  approx- 
imately o°.63  —  "  Sci.  Memoirs.  The  above  table  makes  apparent  the 
real  meaning  of  Raoult's  Law.  It  also  suggests,  that  inasmuch  as  the 
molecular-lowering  of  water,  which  in  a  number  of  cases  is  as  great  as 
47  and  is  an  exception  to  the  law,  that  its  molecular  weight  is  probably 
greater  than  18.  Ramsay  and  Shields,  from  a  study  of  the  surface-tension 
of  water,  reached  the  conclusion  that  .the  molecular  formula  for  water  in 
the  liquid  state  should  be  (H2O)4  or  H8O4.  This  conclusion  is  confirmed 
by  the  freezing  point  lowerings  of  water  produced  by  dissolved  substances. 

K 

The  value  2°  .61,   (  —  ), 
M 

indicates  that  the  weight  of  solvent  used,  contained  only  one-fourth  of  the 
number  of  molecules  required,  or  in  other  words,  each  molecule  weighed 
four  times  as  much  as  calculated.  Assuming  that  the  weight  of  the  mole- 
cule should  be  18  x  4  =  72,  and  inserting  this  value  in  the  equation, 

K  47 

—  =  (o°.63,  Aver.),  we  have  —  —  =  o°.65. 

M  18x4 

This  value  is  very  close  to  the  other  values  of  the  table. 

It  should  be  remembered  that  Raoult's  Law  is  only  an  approximation, 
but  one  of  much  value  in  correlating  many  facts. 

Experiment  XVIII.  —  Lowering  of  the  Solution  Tension. 

Half  fill  a  250  cm.3  or  500  cm.3  flask  with  distilled  H2O;  (The  flask 
should  have  a  long  narrow  neck)  saturate  the  water  with  ether;  add  H2O 
and  ether  repeatedly  until  flask  contains  a  saturated  aqueous  solution  of 
ether  which  rises  into  the  neck  of  the  flask  for  a  distance  of  about  2  cm.3; 
there  should  be  a  layer  of  ether  about  i  cm.  in  thickness  above  the  aqueous 
solution.  The  height  of  each  should  be  marked  with  pieces  of  gummed 


102  EXPERIMENTAL  CHEMISTRY. 

labels.  Now  add  3  or  4  cm.3  of  benzol,  or  toluol,  either  of  which  is  quite 
soluble  in  ether;  mark  its  height  on  neck  of  flask;  shake  vigorously  for  a 
minute  (avoid  spilling  the  contents  of  the  flask) ;  allow  flask  to  stand  un- 
disturbed for  a  few  minutes.  Observe  the  height  of  the  aqueous  solution. 
Note  the  thickness  of  the  layer  of  ether.  Is  there  more  or  less  ether  in 
solution  in  the  water  than  there  was  before  adding  the  benzol  to  the  ether  ? 
Why  did  the  layer  of  ether  increase  in  thickness  by  the  process  ? 

The  answer  to  the  question  as  to  whether  the  process  of  solution  is  to 
be  regarded  as  purely  chemical  or  physical  must  be  delayed  until  such  time 
as  the  evidence  sh^.ll  preponderate  for  one  or  the  other  of  the  two  views, 
which  is  certainly  not  the  case  with  oar  present  information.  The  question 
has  been  much  discussed,,  am}  jbecause  of  the  alteration  in  value  of  well 
defined  properties  of  a  sutstance  when  it  is  dissolved  in  a  liquid,  there  are 
many  chemists  who  regard  solution  as  a  chemical  process.  As  a  rule, 
the  sum  of  the  values  of  a  property  common  to  solute  and  solvent  is  not 
the  value  of  this  property  for  the  solution.  Thermal  and  volume  rela- 
tions, solutions  of  constant  boiling  point,  hydration  and  other  phenomena 
which  appear  during  the  dissolving,  point  to  the  existence  of  a  certain 
affinity  between  solute  and  solvent.  There  is  little  doubt  but  that  the 
two  react  and  influence  each  other's  properties,  but  as  to  the  origin  and 
nature  of  this  influence  chemists  are  unable  to  offer  a  probable  expla- 
nation. 

Those  chemists  who  maintain  that  solution  should  not  be  regarded  as  a 
chemical  process  point  to  the  conduct  of  the  solute  in  dilute  solutions,  with 
respect  to  its  volume,  temperature  and  pressure  relations,  i.e.,  the  appli- 
cability of  the  gas  law.  In  such  cases  the  influence  of  the  solvent  may  be 
neglected  altogether.  Again,  we  find  that  the  quantity  of  a  solid  which 
a  liquid  will  dissolve  varies  with  temperature;  such  variation  in  the  com- 
position of  chemical  compounds  is  unknown.  The  impossibility  of  ex- 
pressing the  concentration  of  saturated  solutions  in  terms  of  integral  multi- 
ples of  the  chemical  combining  weights,  and  the  easy  recovery  of  the  solid 
solute  by  evaporation,  together  with  the  above  mentioned  phenomena  tend 
to  suggest  that  the  process  of  dissolving  is  physical. 

The  subject  of  solutions  will  be  discussed  in  a  future  chapter. 

PROBLEMS. 

i. — How  many  grm.  of  KOH  will  be  required  to  prepare  500  cm.3  of  a 
i  N  solution?  An  .01  N  solution?  A  5  N  solution? 

2. — Calculate  the  relative  number  of  grm.  of  each  solute  in  normal  solu- 
tions of  H2SO4,  HC1,  H3PO4.  In  3  N  solutions.  In  2  N  solutions. 

3. — How  many  grm.  of  HNO3  in  1000  cm.  of  a  5  per  cent,  solution 
(aqueous)  ?  If  10  grm.  of  NaOH  are  dissolved  in  sufficient  H2O  to  give 
a  5  per  cent,  solution  of  the  solute,  what  was  the  total  volume  of  the  solu- 
tion? 

4. — How  many  cm.3  of  water  will  have  to  be  added  to  10  grm.  of  H2SO4 
(sp.  gr.  1.84)  to  yield  a  10  per  cent,  solution  of  acid? 


SOLUTIONS.  103 

5. — If  i  1.  of  H2O  absorbs  i  1.  of  CO2  at  o°,  760  mm.  how  many  grams 
of  CO2  gas  are  contained  in  a  bottle  of  carbonic  water  holding  200  cm.s 
of  solution,  the  pressure  being  5  atmospheres  ? 

6. — The  approximate  composition  of  air  is  30.9  percent,  of  O,  and  79.1 
of  N,  by  volume.  At  15°  water  absorbs  0.0299  volumes  of  O  and  0.0148 
of  N,  the  pressure  of  each  being  that  of  the  atmosphere.  What  is  the 
composition  of  air  absorbed  in  H2O  ? 

Ans.     By  volume,  34.8  per  cent,  of  O  and  65.2  per  cent,  of  N. 

7. — The  increase  in  the  boiling-point  of  54.65  grm.  of  CS2  caused  by 
the  addition  of  1.4475  grm-  of  P  is  o°.486.  What  is  the  molecular  weight 
of  P  in  CS2?  What  is  the  molecular  formula  if  the  atomic  weight  is  31  ? 

8. — The  molecular  weight  of  a  substance  is  60.  If  10  grm.  of  this 
substance  is  dissolved  in  100  grm.  of  a  solvent,  the  increase  in  the  boiling 
point  is  o°.87.  Calculate  the  molecular  increase  in  the  boiling  point. 

9. — Find  the  molecular  weight  of  oxime,  (CH3)2  CNOH,  if  0.284  grm. 
of  it  causes  a  decrease  of  o°.i55  in  the  freezing  point  of  100  grams  of 
glacial  acetic  acid. 

10. — A  student  working  in  this  laboratory  tabulated  the  following 
data  while  determining  the  molecular  weight  of  turpentine  by  the  "  boiling 
point  method."  Boiling  point  of  pure  solvent  (ether)  as  indicated  by  a 
Beckmann  thermometer,  3°. 25;  weight  of  solvent  18.45  grm-5  weight  of 
turpentine,  0.865  grm-;  boiling  point  of  solution,  3°-96.  Calculate  the 
molecular  weight  of  ether. 

n. — The  following  data  was  recorded  while  determining  the  molecular 
weight  of  sugar  by  "freezing  point  method." 

Freezing  point  of  pure  H2O,  (Beckmann)  5°.475, 

Freezing  point  of  solution  after  adding  i  gram  of  sugar  to  20.09 
grams  of  H2O,  5°.  198.  Calculate  the  molecular  weight  of  sugar. 


CHAPTER  XII. 
ACIDS,  BASES  AND  SALTS. 

NOMENCLATURE. 
CHLORINE. 

Symbol,  Cl.  At.  Wt.  35.45. 

Chlorine  is  an  elementary  substance  belonging  to  the  class  of  "  non- 
metals."  Under  ordinary  laboratory  conditions  it  is  a  heavy,  greenish- 
yellow  gas  which  possesses  an  irritating  odor  and  poisonous  properties. 
It  was  discovered  by  Scheele  in  1774,  who  prepared  it  by  heating  hydro- 
chloric acid  with  the  oxide  of  manganese.  Davy,  however,  proved  its  ele- 
mentary nature  and  gave  it  its  name  in  1810.  It  possesses  a  great  deal  of 
chemical  energy,  i.e.,  shows  great  tendency  to  react  chemically  with  other 
substances,  therefore,  it  is  not  found  in  nature  as  jree  chlorine.  Observe 
during  the  performance  of  the  following  experiment  whether  chlorine 
possesses  the  physical  properties  characteristic  of  metals. 

Experiment  I. — Preparation  of  Chlorine. 

Arrange  a  300  cm.3  flask  (Fig.  22)  provided  with  funnel  tube  and 
a  delivery  tube,  so  that  it  connects  with  washing  bottle  which  should 
be  about  one-fourth  full  of  concentrated  sulphuric  acid.  The  flask 
should  be  placed  on  a  piece  of  iron  gauze  resting  on  a  ring  clamp  or  tripod 
so  that  flask  may  be  heated.  Introduce  into  flask  20  grams  of  manganese 
dioxide  (MnO2),  preferably  in  the  powdered  condition.  Place  six  clean 
dry  bottles  of  about  150  cm.3  capacity  on  your  desk  and  cover  each  with 
a  glass  plate  or  a  sheet  of  heavy  damp  paper;  procure  from  the  assistant 
a  little  powdered  antimony  on  a  watch  glass,  a  strip  of  bright  colored  calico, 
some  litmus  paper,  write  your  name  with  ink  upon  a  piece  of  paper,  4  or  5 
cm.3  of  turpentine  in  a  test  tube  and  a  test  tube  half  filled  with  distilled 
water;  then  arrange  your  hydrogen  generator  so  that  you  can  provide 
a  hydrogen  flame  on  short  notice.  Now  add  by  means  of  funnel  tube, 
sufficient  concentrated  hydrochloric  acid  to  cover  the  MnO2  in  flask;  be 
sure  that  lower  end  of  funnel  tube  is  dipping  into  acid;  heat  flask  very 
gently;  by  means  of  a  delivery  tube  bent  at  right  angles  and  attached 
to  the  wash  bottle  so  that  it  can  be  inserted  into  the  mouth  of  a  bottle, 
collect  gas  by  vertical  displacement  of  air.  When  bottle  presents  a 
greenish-yellow  appearance  remove  it  and  cover  its  mouth  with  glass  plate 
or  damp  paper.  Repeat  operation  until  six  bottles  have  been  filled  with 
the  gas  and  water  in  test  tube  has  been  saturated  with  chlorine.  Then 

IQ\ 


ACIDS,    BASES    AND    SALTS.  105 

proceed  with  the  following  experiments.     The  equation  for  the  above 
reaction  may  be  written  as  follows: 

MnO2  +  4  HC1  —  MnCl4  +  C12  +  2H2O. 
MnCl4  —  MnCl2  +  C12. 

Is  chlorine  lighter  or  heavier  than  air?  If  a  liter  of  hydrogen  weighs 
.0899  grams  under  standard  conditions  what  will  the  same  volume  of 
chlorine  gas  wreigh  under  identical  conditions  ? 

Experiment  II. — Properties  of  Chlorine. 

(a)  Bleaching  properties.  Into  a  bottle  of  gas  insert  pieces  of  dry 
calico,  litmus  paper  and  paper  with  carmine  ink  upon  it;  set  it  aside  for 
fifteen  or  twenty  minutes.  Repeat  above  operation  with  another  bottle 
of  chlorine,  but  moisten  the  articles  before  inserting  them.  Examine  be- 


FIG.  22. 

fore  leaving  laboratory.  Results?  Compare  with  result  when  articles 
were  perfectly  dry.  What  conclusion  would  you  draw?  It  should  be 
observed  that  organic  material  is  the  basis  of  the  coloring  matter  present 
in  above  specimens. 

(b)  Combination  of  antimony  and  chlorine.     Into  a  bottle  of  the   gas, 
drop  a  pinch  of  powdered  antimony.     Result?     Place  another  portion 
of  the  powder  in  a  test  tube  and  heat  it  until  quite  warm,  then  quickly 
introduce  it  into  the  same  bottle  of  gas.     Result  ?     What  factor  induced 
the  reaction  ?     The  valency  of  antimony  is  three,  that  of  chlorine  is  one. 
Write  the  equation  for  reaction. 

(c)  Affinity  of  chlorine  and  hydrogen. 

Insert  "hydrogen  flame"  into  a  bottle  of  chlorine  (Fig.  23).     Result? 


106  EXPERIMENTAL  CHEMISTRY. 

Hold  a  piece  of  moist  blue  litmus  paper  just  outside  of  mouth  of  bottle 
while  "  flame"  is  present  on  inside.  Result  ?  Hold  another  piece  of  blue 
litmus  paper  near  the  mouth  of  the  reagent  bottle  labeled  "  hydrochloric 
acid."  Results?  What  are  your  conclusions  as  to  the  gas  formed  by 
the  burning  of  chlorine  in  hydrogen  ?  Write  equation  for  reaction. 

Was  light  and  heat  evolved  as  the  result  of  the  reaction  ?  What  was 
probably  the  main  source  of  this  energy?  Was  it  an  endo-  or  exother- 
mic reaction?  Interpret  following  equation: 

H,  (i  gram)  +  Cl,  (35.45  grams)  — >  HC1,  (36.45  grams)  +  22,000  cals. 
(d)  Chemical  action  on  turpentine  (C10H16). 

Place  a  strip  of  filter  paper  in  the  test  tube  containing  the  C10  H16  and 
heat  tube  gently  until  it  is  warm.     Turpentine  is  very  inflammable.     In- 
troduce paper  saturated  with  warm  turpen- 
tine into  a  bottle  of  chlorine  gas.     Result  ? 
If  you  were  told  that  the  black  appear- 
ance of  paper  was  due  to  free  carbon  (soot) 
what  would  be  your  conclusion  ?   Equation  ? 
(e)  Substituting  power  of  chlorine, 
i.  Make  a  few  strips  of  "test  paper"  by 
dipping  pieces  of  filter  paper  in  a  solution 
of    potassium    iodide    (KI)    with   which  a 
clear  solution  of  starch   has  been   mixed. 
Free  iodine  will  impart  a  blue  or  bluish- 
FIG.  23.  black  color  to  starch.     Upon  a  strip  of  test 

paper  pour  a  few  drops  of  the  water  saturated 
with  chlorine.     Results?     Conclusions?     Equation? 

2.  Pass  hydrogen  sulphide  (H2S)  through  a  water  solution  of  chlorine 
and  observe  effect.     Before  passing  the  H2S  into  the  solution  notice  the 
action  of  the  latter  upon  litmus  paper,  both  red  and  blue.     Compare  the 
result  with  a  similar  test  made  after  the  H2S  has  been  passed  into  it.     In 
view  of  observed  data,  how  would  you  explain  results  ?     Equation  ? 

3.  Optional.      (L.  T.)      Fill  a  test  tube  with  a  saturated  solution  of 
chlorine;  invert  it — mouth  under  water;  place  it  where  direct  sunlight 
may  fall  upon  the  tube;  examine  after  a  day  or  two.     Results?     Test 
water  of  tube  with  litmus  paper.     Results?     If  any  gas  has  accumu- 
lated in  upper  portion  of  tube,  test  it  with  a  glowing  splinter.     Results  ? 
What  are  your  conclusions  as  to  reaction  ?     Equation  ? 

4.  Optional.     Introduce  into  a  bottle  of  chlorine  a  piece  of  metallic 
sodium  heated  gently  in  a  dephlegrating  spoon.     Result?      Equation? 

Experiment   III.      Optional    Methods    for    Preparation    of   Chlorine. 
(Oxidation  Processes.) 

(a)  Use  equal  parts  by  weight  of  NaCl  and  MnO2;  add  strong  sulphu- 
ric acid. 

Mn02  +  NaCl  +  tt^O,  —  MnSO4  +  Na,SO4  +  2H2O  +  2CL 


ACIDS,  BASES  AND  SALTS.  107 

For  stages  of  reaction  (Instruction). 

(b)  Potassium  permanganate,   with   hydrochloric   acid  added  slowly 
from  a  "dropping"  funnel. 

KMnO4  +  8HC1  -»  KC1  +  MnCl2  +  4H2O  +  sCl. 

(c)  To  a  few  crystals  of  potassium  chlorate,  add  concentrated  hydro- 
chloric acid.     (This  is  a  convenient  and  efficient  method  for  preparing 
small  quantities  of  the  gas.     It  is  frequently  resorted  to  by  analytical 
chemists.)     The  reaction  may  be  represented  by  the  following  equation: 


2KC1O3  +  4HC1  —  2KC1  +  C12O4  +  2H3O  +  2C1. 

It  must  be  evident  to  the  student  that  chlorine  possesses  considerable 
chemical  energy,  and  that  it  manifests  a  remarkable  affinity  for  hydrogen 
— the  majority  of  the  foregoing  phenomena  being  explained  by  this  fact. 
Its  bleaching  properties  in  the  presence  of  moisture  appear  to  be  due  to 
the  reaction  whereby  the  hydrogen  is  withdrawn  from  the  water,  liberat- 
ing nascent  oxygen  which  oxidizes  the  coloring  compounds,  thus  form- 
ing new  substances  which  are  colorless.  Chlorine  has  frequently  been 
called  an  "oxidizer"  on  account  of  this  type  of  reaction.  Frequently 
compounds  are  " broken  up,"  i.e.,  suffer  decomposition  because  of  the 
withdrawal  of  hydrogen  by  chlorine. 

That  which  is  of  dominant  interest  to  us  at  this  time  is  the  compound 
which  is  formed  as  the  result  of  the  combination  of  hydrogen  and  chlo- 
rine, and  which  in  the  light  of  former  rules  of  nomenclature,  might  be 
termed  hydrogen  chloride.  The  properties  of  this  latter  substance  re- 
sembled those  of  hydrochloric  acid.  This  leads  us  to  a  study  of  a  class 
of  compounds  called  "acids." 

ACIDS. 

In  the  study  of  the  element  chlorine  which  is  a  typical  non-metal, 
a  compound  containing  chlorine  and  hydrogen  was  prepared,  which  in  a 
number  of  its  properties  resembled  a  class  of  compounds  called  acids. 
The  investigation  of  the  properties  of  this  compound  may  warrant  the 
making  of  certain  generalizations  which  will  be  of  service  to  us. 

Experiment  I  (L.T.). — Electrolysis  or  Analysis  of  Hydrochloric  Acid. 

Fill  a  Hoffman  V-tube  for  the  electrolysis  of  hydrochloric  acid  with  a 
dilute  solution  of  the  shelf  reagent.  Connect  the  positive  pole  of  the 
battery  with  the  end  of  the  V-tube  which  is  not  ordinarily  sealed,  but  is 
closed  with  a  cork  through  which  a  platinum  electrode  passes.  Pass  the 
current;  hold  a  piece  of  "test  paper"  near  the  open  end  of  tube  for  three 
or  four  minutes.  Results  ?  Conclusions  ? 

Test  the  gas  collecting  in  the  other  arm  of  the  tube  by  bringing  a  lighted 
match  near  it.  Results  ?  Identity  of  gas  ?  What  are  your  conclusions 
as  to  the  composition  of  hydrochloric  acid  ?  Is  your  view  supported  by 
any  of  the  experiments  performed  with  chlorine?  Explain.  At  which 


io8 


EXPERIMENTAL  CHEMISTRY. 


electrode  was  free  chlorine  (non-metal)  liberated  ?     This  is  a  characteristic 
of  non-metals. 

Experiment  II. — Volumetric  Composition  of  Hydrochloric  Acid. 
(a)  Fill  the  Hoffman  apparatus  which  was  used  for  the  electrolysis  of 
water,  with  a  solution  of  hydrochloric  acid.     Pass  the  current  for  fifteen 
or  twenty  minutes  simply  to  saturate  the  water  with  hydrogen  and  chlo- 
rine gases,  as  the  latter  is  particularly  soluble;  open  stop  cocks,  forcing 
out  the  gases;  now  proceed  with  the  electrolysis  to  determine 
the  ratio   of   the   combining  volumes   of   hydrogen   and 
chlorine.     Recalling  Avogadro's  rule,  what  would  you  say 
in  regard  to  the  relative  number  of  atoms  of  each  gas  in  a 
molecule  of  the  acid? 

(b)  Optional.       Relative    volumes    of    hydrogen    and 
chlorine  in  hydrogen  chloride. 

Generate  hydrogen  chloride  as  suggested  in  above  ex- 
periments, then  dry  gas  by  bubbling  it  through  a  wrash 
bottle  filled  with  sulphuric  acid,  or  generate  it  by  allowing 
concentrated  sulphuric  acid  to  drop  slowly  from  a  separa- 
tory  funnel  into  a  flask  containing  concentrated  hydro- 
chloric acid.  Fill  a  colorimetric  tube  or  eudiometer  (Fig. 
24)  with  the  dry  gas.  Introduce  about  10  grams  of  sodium 
amalgam  and  quickly  cork  the  open  end.  Raise  and  lowrer 
either  end  of  the  tube  so  that  the  amalgam  may  come  into 
contact  with  all  portions  of  the  gas.  Uncork  the  tube 
under  mercury.  Result  ?  Adjust,  so  that  liquid  on  out- 
side and  inside  of  tube  are  at  same  level.  What  is  the 
volume  ?  Place  thumb  over  open  end  of  tube  and  lift  it 
from  the  liquid;  invert;  remove  thumb  and  apply  a  lighted 
FIG.  24.  match.  Results?  What  gas  was  it? 

What  are  your  conclusions  as  to  the  relative  volumes  of 
the  two  constituents  of  hydrogen  chloride  ?  As  to  the  volume  of  either 
constituent  relative  to  the  total  volume  ? 

Experiment  III. — Preparation  and  Properties  of  Hydrogen  Chloride  Gas. 

(a)  Place  a  small  quantity  of  ammonium  chloride  in  a  test  tube;  add 
a  few  drops  of  concentrated  sulphuric  acid;  hold  a  piece  of  blue  litmus 
near  mouth  of  tube.  Results  ?  Repeat  with  red  litmus  paper.  Results  ? 
Blow  your  breath  across  mouth  of  tube.  Results?  Pour  a  few  cm.3 
of  ammonium  hydroxide  into  a  test  tube,  then  bring  its  mouth  near  to 
mouth  of  tube  in  which  gas  is  being  generated.  Results? 

2NH4C1  +  H2SO4  —  (NH4)2SO4  -f  THC1. 

1  mol.  grm.  wt.  -2  mol.  grtn.  wts. 

2  mol.  grm.  wts.  i  mol.  grm.  wt. 

To  what  class  of  chemical  reactions  does  the  above  reaction  belong? 
Write  equation  for  the  reaction  that  occurred  when  the  two  test  tubes 
with  contents  were  brought  close  together. 


ACIDS,  BASES  AND  SALTS.  1 09 

Experiment  IV. — Study  of  the  Properties  of  the  Gas  Evolved  from  Hy- 
drochloric Acid. 

Now  half  fill  a  test  tube  with  the  shelf  reagent  known  as  "  hydrochloric 
acid."  Using  this,  repeat  the  above  experiment  in  its  various  parts,  and 
record  results. 

Does  there  seem  to  be  any  similarity  in  the  properties  of  hydrogen 
chloride  gas  and  the  gas  evolved  from  hydrochloric  acid  ? 

Experiment  V. — Laboratory  Method  for  Preparation  of  Hydrogen  Chlor- 
ide Gas. 

Arrange  apparatus  as  you  did  for  the  preparation  of  chlorine  with  this 
exception,  that  you  raise  the  tube  of  the  wash  bottle  which  ordinarily 
dips  into  the  washing  fluid  (water  in  this  experiment)  so  that  its  lower  end 
just  escapes  the  surface  of  the  water.  Pour  15  cm. 3  of  water  into  a  beaker; 
add  slowly  and  while  constantly  stirring  35  cm.3  of  concentrated  sulphuric 
acid;  let  the  mixture  cool.  Introduce  35  grams  of  sodium  chloride  into 
generating  flask  and  observe  that  all  parts  of  system  of  apparatus  are  prop- 
erly arranged.  Place  upon  the  desk  a  f ewr  dry  bottles,  and  a  label  cut  into 
narrow  strips  about  a  cm.  long.  Set  a  dry  bottle  under  delivery  tube  of  wash 
bottle  to  collect  gas  by  vertical  displacement  of  air.  Add  slowly  through 
funnel  tube  about  one-half  of  cooled  mixture  of  water  and  acid, — if  neces- 
sary add  all  of  it,  and  heat  gently.  Fill  three  bottles  with  the  hydrogen 
chloride  gas.  Half  fill  a  clean  test  tube  with  distilled  water;  mark  its 
height  by  means  of  a  piece  of  label;  permit  the  gas  to  bubble  into  water  in 
tube  until  it  is  saturated;  this  test  tube  should  be  placed  in  a  bath  which 
is  at  a  temperature  as  near  o°  C.  as  can  conveniently  be  arranged;  at 
least,  place  tube  in  a  bottle  filled  with  cold  water.  While  waiting  for 
water  to  become  saturated,  proceed  with  following  experiments.  The 
above  reaction  may  be  represented  as  follows: 

NaCl  +  H2SO4  —  NaHSO4  +  HCL 
NaHSO4  +  NaCl  -»  Na^SO,  +  HCL 

Experiment  VI. — Solubility  of  Hydrogen  Chloride  Gas. 

Invert  one  of  the  bottles  of  gas  over  a  pneumatic  trough,  with  its  mouth 
under  water.  Results  ?  Conclusions  ?  Proofs  ? 

Experiment  VII. — Density  of  Hydrogen  Chloride  Gas. 

Using  another  bottle  of  the  gas,  devise  a  method  which  will  prove 
approximately  that  the  gas  is  heavier  than  air.  Can  gas  be  "  poured  ?" 

Experiment  VIII. — Properties  of  Hydrogen  Chloride  Gas  in  Solution 
and  Hydrochloric  Acid. 

Before  proceeding  with  this  experiment,  take  up  Exp.  IX.  Divide 
the  aqueous  solution  of  the  gas  into  three  equal  parts. 

(a)  To  one  part  in  a  test  tube  add  a  piece  of  granulated  zinc.  Results  ? 
Hold  a  lighted  match  near  the  mouth  of  the  test  tube.  Result  ?  Con- 


110  EXPERIMENTAL  CHEMISTRY. 

elusions?     Repeat  this,  and  all  of  the  following  experiments,  by  using 
hydrochloric  acid  from  the  shelf.     Results  ? 

(b)  To  the  second  part  of  the  solution  add  a  little  sodium  carbonate. 
Results  ? 

(c)  Divide  the  third  part  into  four  portions  and  dilute  each  with  twice 
its  volume  of  water.     To  one  of  the  portions  add  a  few  drops  of  silver 
nitrate.     Result?     To  another  portion  add  a  few  drops  of   mercurous 
nitrate.     Results  ?     To  the  third  portion  add  lead  nitrate.     Results  ? 
Divide  the  fourth  portion  into  two  parts.     To  one  part  add  sodium  hy- 
droxide.    Results  ? 

To  the  second  part  add  three  times  its  volume  of  water.  Taste  it  by 
placing  a  drop  or  two  on  the  -tongue  by  means  of  stirring  rod.  Does  it 
possess  a  caustic,  lye-like,  or  an  acid,  sour  taste  ? 

Dip  pieces  of  blue  and  red  litmus  paper  into  the  solution.     Results? 

What  are  your  conclusions  as  to  the  relation  between  hydrogen  chloride 
gas  in  aqueous  solution  and  the  shelf  reagent,  hydrochloric  acid? 

Experiment  IX. — (Quant.)  Optional.  Density  of  a  Solution  of  Hydrogen 
Chloride  Gas  (Hydrochloric  Acid). 

When  liquid  in  above  test  tube  (Exp.  V)  is  apparently  saturated  with  gas, 
remove  it  and  note  volume  of  contents.  Results  ?  Now  carefully  remove 
all  traces  of  label  and  wipe  dry  outside  and  inside  of  tube  above  solution; 
weigh  tube  and  contents.  Record  weight;  mark  height  of  solution  with 
label;  pour  contents  of  tube  into  another  tube  to  be  reserved  for  preceding 
experiment;  wash  tube  and  fill  with  distilled  water  to  top  of  "label;" 
remove  label  and  wipe  tube  as  before;  weigh  empty  tube  when  clean  and 
dry.  Weight  ?  Find  density  of  solution.  Record  all  calculations. 

Experiment  X. — General  Properties  of  Acids. 

Make  very  dilute  solutions  of  the  following  acids,  nitric,  sulphuric,  and 
acetic.  Taste  a  drop  of  each.  Test  them  with  both  blue  and  red  lit- 
mus paper.  Place  a  piece  of  granulated  zinc  in  a  dilute  solution  of  each. 
Test  each  acid  with  a  solution  of  sodium  carbonate.  Make  a  record  of 
all  observations.  Do  all  of  the  acids  have  some  properties  in  common  ? 
Write  the  formulae  for  hydrochloric,  nitric,  sulphuric,  and  acetic  acids. 
What  element  is  possessed  in  common  by  all  acids?  Would  you  sus- 
pect a  causal  relation  between  the  element  and  the  properties  possessed 
in  common  by  acids  ? 

Experiment  XI. — Each  Acid  Has  Characteristic  Properties. 

Procure  from  side  shelf  reagents  4  or  5  cm. 3  of  lead  nitrate  in  a  test 
tube,  and  an  equal  volume  of  barium  chloride  in  another  tube.  Divide 
the  lead  nitrate  into  two  parts.  To  one  part  add  a  few  drops  of  dilute 
hydrochloric  acid;  to  the  other  add  dilute  nitric  acid.  Record  results. 

Repeat  the  above  using  nitric  acid  and  sulphuric  acid  and  the  barium 
chloride  solution.  Record  results.  Would  you  say  that  each  acid  pos- 


ACIDS,    BASES    AND    SALTS.  Ill 

sesses  properties  peculiar  to  itself  in  addition  to  its   general  acid  prop- 
erties ? 

Write  equations  for  all  reactions,  underscoring  the  substance  which 
separated  out  in  the  solid  form  (precipitate).  Is  the  precipitate  in  equilib- 
rium with  molecules  in  solution? 

It  is  difficult  to  accurately  define  an  acid  at  this  stage  of  our  work. 
From  the  foregoing  experiments  it  may  be  deduced  that  acids  are  com- 
pounds which  may  be  either  solids,  liquids,  or  gases.  They  possess 
the  power  to  alter  certain  vegetable  colors,  namely,  to  turn  blue  litmus 
red;  to  act  upon  metals,  which  displace  hydrogen;  they  have  a  sour  taste 
and  usually  contain  a  non-metal  (an  electro-negative  element)  united  with 
hydrogen  or  hydrogen  and  oxygen.  It  is  generally  conceded  that  it  is 
the  replacable  hydrogen  of  these  compounds  when  in  aqueous  solution 
which  give  them  their  general  acid  properties. 

Hydrascids.         (HC1,H2S). 
ACIDS, 

Oxascids.  (HC103,H2S04). 

Organic.  (HC2H3O2,  H2C4H4O6). 

ACIDS, 

Inorganic.  (HNO3,  H3PO4). 

Oxascids  are  sometimes  looked  upon  as  non-metallic  oxides  combined 
with  water.  The  oxide  is  called  the  anhydride  (H2O.SO3,H2O.N2O5). 

The  part  of  an  acid  formula  which  remains  after  the  hydrogen  is  re- 
moved is  sometimes  spoken  of  as  the  acid  radical. 

Hypo-ous.  (HC1O). 

-ous.  (HC102). 
OXASCIDS, 

-ic.  (HC10S). 

per-ic.  (HC104). 

The  above  nomenclature  is  a  type  of  that  which  is  resorted  to  in  order 
to  distinguish  between. a  number  of  compounds  closely  related  in  com- 
position. The  endings  indicate  the  relative  degrees  of  oxidation. 

Binary  Compounds  are  those  which  are  composed  of  two  elements, 
M2O.  The  names  of  such  substances  end  .in  ide.  This  rule  relative  to 
nomenclature  has  preference  over  any  other  rule.  Ternary  Compounds 
are  those  composed  of  three  elements,  as,  H2SO4. 

SODIUM. 

Symbol,  Na.  At.  Wt.  23.05. 

Sodium  is  a  typical  metal.  In  dividing  the  elementary  substances  into 
metals  and  non-metals,  it  should  be  remembered,  and  as  we  shall  see  later, 
the  line  of  demarcation  is  nowhere  distinctly  drawn.  Midway  between 


112  EXPERIMENTAL  CHEMISTRY. 

these  proposed  classes  there  are  such  elements  as  arsenic,  antimony  and 
bismuth,  whose  chemical  and  physical  properties  permit,  under  varying 
conditions,  a  classification  with  either  the  metals  or  non-metals.  These 
elementary  substances  occupying  intermediate  ground  are  sometimes 
called  metalloids.  It  is  better  to  regard  the  elements  as  constituting  a 
series  with  a  regular  gradation  of  properties. 

Sir  Humphry  Davy  succeeded  in  preparing  small  quantities  of  sodium 
by  electrolysis  of  fused  sodium  hydroxide,  about  1807. 

Experiment  I. — Properties  of  Sodium. 

(Metallic  sodium  is  usually  kept  under  kerosene.  It  should  always  be 
handled  by  means  of  dry  forceps.) 

(a)  Place  a  piece  of  sodium  upon  a  dry  paper  and  cut  off  a  piece  the 
size  of  a  small  pea.     Is  the  sodium  hard?     Observe  the  color,  luster, 
and  the  effect  of  air.     Half  fill  the  pneumatic  trough  with  water  which 
does  not  turn  red  litmus  paper  to  blue.     Drop  the  small  piece  of  metal 
upon  the  water  and  step  back  from  the  trough.     Does  the  metal  float? 
Is  there  evidence  of  chemical  action  taking  place  ?    Fasten  a  match  to  a 
rod;  light  match,  and  apply  flame  to  piece  of  metal.     Is  there  evidence 
of  a  flame  around  the  sodium?     A  flame  implies  the  existence  of  a  gas. 
Where  does  the  gas  come  from? 

(b)  Wrap  a  piece  of  sodium  in  a  little  paper;  place  it  in  a  wire  gauze 
basket;  hold  it  under  water  and  collect  the  gas  which  escapes  in  a  test  tube 
by  displacement  of  water;  apply  a  lighted  match  to  mouth  of  tube. 
Results  ?     What  gas  ?     The  sodium  has  apparently  displaced  what  from 
the  molecule  of  water? 

(c)  Test  the  water  in  the  trough  with  red  litmus  paper.     Result? 
Take  a  little  of  the  water  between  the  fingers. 

Pour  3  cm.3  of  the  shelf  reagent  labeled  sodium  hydroxide  into  a  test  tube 
and  dilute  with  an  equal  volume  of  water.  Repeat  above  tests.  Would 
you  say  that  they  are  identical  substances  ?  If  so  how  much  of  the  hydro- 
gen was  displaced  from  the  molecule  of  wrater  ?  Write  equation  to  repre- 
sent the  reaction  of  sodium  and  water.  What  is  the  valency  of  sodium  ? 

(d)  Recall  the  Exp.  in  which  metallic  sodium  was  introduced  into 
chlorine.     Write  equation. 

(e)  Flame  color  of  sodium  compounds.     Dip  a  platinum  wire  into  a 
little  hydrochloric  acid;  hold  it  in  flame  until  it  imparts  no  color  to  the 
latter;  dip  wire  into  a  little  sodium  hydroxide  and  hold  wire  in  flame. 
Is  the  flame  colored  ?     What  is  the  color  ?     Repeat  all  of  above  operations 
using  the  water  of  the  pneumatic  trough.     Results?     Conclusions? 

Sodium  is  a  silver-white  metal;  it  loses  its  metallic  lustre  on  exposure  to 
damp  air  due  to  the  fact  that  under  such  conditions  it  readily  oxidizes; 
it  is  quite  soft  at  laboratory  temperatures,  but  hard  at  — 20°  C.;  melting 
point,  96.5°  C.;  specific  gravity,  .97;  conductor  of  heat  and  electricity;  is 
an  electro-positive  element,  separating  out  at  negative  electrode,  but  imme- 
diately reacts  on  the  water  forming  a  compound  known  as  sodium  hy- 


ACIDS,  BASES  AND  SALTS.  113 

droxide  and  liberating  hydrogen  instead;  it  possesses  relatively  a  great 
amount  of  chemical  energy,  combining  so  readily  with  substances  that 
it  is  not  found  in  the  free  state  in  nature.  We  have  seen  it  displace  an 
hydrogen  atom  from  the  molecule  of  water  and  combine  with  the  hydroxyl 
(OH),  yielding  a  compound  known  generically  as  a  base  or  an  alkali,  and 
specifically,  as  sodium  hydroxide  or  sodium  hydrate,  which  has  the  property 
of  turning  red  litmus  paper  blue. 

BASES. 

Experiment  I. — General  Properties  of  Bases. 

Place  3  cm.3  of  each  of  the  following  in  different  test  tubes,  and  dilute 
each  with  four  times  its  volume  of  water: — sodium  hydroxide,  potassium 
hydroxide,  ammonium  hydroxide  and  calcium  hydroxide.  Test  each  by 
means  of  red  and  blue  litmus  paper,  tasting  a  drop,  and  by  taking  a  little 
between  the  fingers.  Tabulate  data.  What  are  your  conclusions  as  to 
their  similarity  with  respect  to  the  properties  of  changing  red  litmus  to 
blue  litmus,  alkaline  taste  (alkalinity,  causticity),  and  lye-like  or  soapy 
feeling  to  touch  ? 

Write  the  formula  of  each.  Is  the  hydroxyl  common  to  all  ?  Would 
you  infer  a  causal  relation  between  properties  in  common  and  constituents 
in  common  ?  Ask  the  instructor  to  show  solid  NaOH  and  KOH,  to  you. 

Experiment  II. — Each  Base  Has  Characteristic  Properties. 

To  a  few  cm.3  of  lead  nitrate  solution  add  a  few  cm.3  of  dilute  hydro- 
chloric acid.  Divide  the  precipitate  into  two  parts.  To  one  part  add 
considerable  NaOH  (reagent).  Result?  To  the  other  portion,  add 
NH4OH  in  excess.  Result? 

Using  a  solution  of  silver  nitrate  repeat  above  experiments.  Results  ? 
Write  equations  for  above  reactions. 

As  the  result  of  the  combination  of  sodium  and  the  hydroxyl  (OH) 
we  have  a  compound  which  is  typical  of  a  large  group  of  compounds  which 
are  known  collectively  as  bases.  They  seem  to  possess  properties  which 
are  the  very  opposite  of  those  possessed  by  acids.  Bases  have  the 
power  to  change  red  litmus  to  blue;  they  have  an  alkaline  or  bitter  taste 
and  feel  soapy  to  the  touch.  The  term  base  has  been  very  loosely,  and 
frequently,  inaccurately  applied.  It  is  now  usually  restricted  to  the 
hydroxides  or  hydrates  of  the  metals.  It  seems  to  be  quite  well  established 
that  the  basic  or  alkaline  properties  of  the  bases  in  aqueous  solution,  are 
due  to  the  action  of  the  hydroxyl.  Our  concept  of  a  base,  in  regard  to  its 
constituents  is,  that  it  is  a  metallic  element  (metal)  combined  with  one  or 
more  hydroxyls,  the  number  depending  on  the  valency  of  the  metal.  As 
would  be  expected  the  above  enumerated  properties  are  possessed  in 
varying  degrees  of  intensity  by  the  different  bases. 

Strong  bases,  like  sodium  hydroxide  and  potassium  hydroxide,  are  fre- 
8 


114  EXPERIMENTAL  CHEMISTRY. 

quently  called  alkalies.     Again,  sodium  hydroxide  and  potassium  hy- 
droxide are  respectively,  spoken  of,  as  caustic  soda  and  caustic  potash 

Vegetable.  Fixed. 

Alkalies  Alkalies 

Mineral.  Volatile. 

The  bases  are  commonly  distinguished  from  one  another  by  using  the 
name  of  the  metal  before  the  word  hydroxide. 

Bases  are  not  infrequently  considered  as  being  composed  of  a  metallic 
oxide  and  water.  The  oxide  is  called  a  basic  oxide.  BaO  +H2O — » 
Ba  (OH)2. 

The  question  might  logically  be  raised  at  this  time,  what  will  occur  if 
an  acid  and  a  base  (their  properties  are  the  opposite  of  one  another)  are 
brought  together? 

NEUTRALIZATION — SALTS. 

Experiment  I. — Electrolysis  of  a  Salt. 

Make  a  strong  solution  of  sodium  sulphate.  Test  the  solution  with 
both  red  and  blue  litmus  paper.  Is  the  solution  neutral  ?  Place  solution 
in  U-tube  for  electrolysis  after  having  added  a  sufficient  quantity  of  lit- 
mus solution  to  give  a  decided  blue  color  to  the  entire  volume  of  liquid. 
What  changes  of  color  occur  when  current  is  passed  ?  What  is  liberated 
at  either  pole  ?  Write  equations  to  represent  all  reactions. 

Experiment  II. — A  Quantitative  Study  of  the  Interaction  of  Acids  and 
Bases. 

Assemble  two  burettes  (Fig.  25)  and  clean  them  thoroughly.  Clamp 
them  into  proper  position.  Fill  the  burette  at  your  right  with  dilute  hydro- 
chloric acid  (i  of  acid  to  20  of  water).  Fill  the  left  burette  with  a  dilute 
solution  of  sodium  hydroxide  (i  part  of  shelf  reagent  to  10  of  water). 
Run  out  acid  and  alkali  so  that  both  burettes  give  zero  reading.  Clean 
a  small  Erlenmeyer  flask;  place  a  piece  of  white  paper  under  either 
burette.  Place  flask  under  "  acid"  burette  and  run  into  it  9  cm.3  of  acid. 
Be  sure  to  read  from  the  lower  side  of  meniscus.  Add  a  few  drops  of 
litmus  solution  or  a  small  piece  of  red  litmus  paper.  Now  place  flask 
under  "alkali"  burette  and  introduce  cautiously  enough  of  the  alkali  to 
just  turn  the  litmus  color  to  a  permanent  blue.  From  time  to  time  shake 
the  flask  vigorously  to  bring  alkali  and  acid  into  intimate  contact,  or  stir 
writh  a  glass  rod,  but  don't  take  the  rod  out  of  the  flask.  Draw  from  the 
burettes  alternately  as  is  necessary  to  bring  the  solution  to  the  point  when  a 
drop  of  either  will  cause  a  change  in  color.  When  this  point  has  been 
reached  the  solution  is  neutral.  Record  the  number  of  cm.3  of  each  used. 
The  process  is  known  as  neutralization.  Taste  the  solution.  Place  15  or 
20  cm.3  in  an  evaporating  dish  and  evaporate  to  dryness.  Taste  the  salt. 
Has  it  a  familiar  taste  ?  While  waiting  for  the  solution  to  evaporate,  re- 
peat first  part  of  operation  by  finding  how  many  cm.3  of  the  alkali  will  be 


ACIDS,  BASES  AND  SALTS.  115 

required  to  neutralize  15  cm.s  of  acid;  then  repeat  again,  using  12  cm.3  of 
acid.     Tabulate  all  data. 

Does  the  Law  of  Definite  Proportions  receive  a  verification  by  your 
results  ? 


FIG.  25. 

If  time  permits,  other  acids  and  bases  may  be  used  and  their  ratios 
determined.  Write  equations. 

What  do  you  have  in  the  acid  that  you  do  not  have  in  the  neutral  salt  ? 
Answer  the  same  question  in  regard  to  the  alkali.  What  has  become  of 
these  two  substances  ?  Is  it  possible  that  their  removal  from  the  neutral 
substances  is  in  any  way  related  to  the  property  of  "neutrality?" 


1 1 6  EXPERIMENTAL  CHEMISTRY. 

Write  the  equations  representing  the  respective  reactions  of  H2SO4, 
HC1,  and  HN03  with  KOH,  NH4OH,  and  Ca(OH)2. 

Experiment  III. — Interaction  of  a  Metallic  Oxide  and  an  Acid. 
ZnO  +  2HC1  —  ZnCl2  +  H2O. 

Experiment  IV. — Interaction  of  Metals  and  Acids. 
Zn  +  H2S04-*ZnS04  +  H,. 

Experiment  V. — Action  of  an  Hydroxide  as  a  Base  and  as  an  Acid. 

To  a  few  cm.3  of  Pb  (NO3)2  add  a  few  cm.3  of  NaOH.  Note  the  pre- 
cipitate. Filter.  Divide  the  precipitate  into  two  parts.  To  one  part 
add  HNO3;  to  the  other,  add  an  excess  of  NaOH.  Is  there  any  similarity 
in  the  action  of  the  acid  and  the  base  in  these  latter  reactions  ?  Write 
equations. 

Experiment  VI. — Action  of  a  Base  on  a  Salt. 

To  3  cm.3  of  Fe  SO4  add  a  few  cm.3  of  NH4OH.  Results?  Equa- 
tion? 

Experiment  VII. — Action  of  an  Acidic  Oxide  upon  a  Base. 

CO2  +  Ca  (OH)2-*CaCO3  +  H2O. 
Experiment  VIII. — Interaction  of  Salts. 

HgNO3  +  Na   Cl  ->  HgCl  +  NaNO3. 

Experiment  IX. — Preparation  of  a  Basic  Salt. 

Recall  the  experiment  in  which  BiCl3  which  is  soluble,  was  changed  into 
an  insoluble  basic  salt  by  hydrolysis.  Define  Hydrolysis. 

H20  ->  H-  +  OH' 
BiCls  +  2H-  +  2OH'  —  Bi(OH)2Cl  +  2HC1 

1 
BiOCl  +  1^0. 

Experiment  X. — Preparation  of  an  Acidic  Salt. 

Fill  a  burette  with  NaOH  as  found  in  the  reagent  bottles.  Fill  another 
burette  with  a  cold  solution  of  H2SO4  (i  of  acid  to  2  of  H2O).  Neutralize 
5  cm.3  of  acid  with  NaOH.  Record  number  of  cm.3  used.  Evaporate 
solution  until  it  becomes  saturated.  Crystals  will  separate  out  when  it 
becomes  cold.  To  the  same  volume  of  NaOH  as  used  above,  add  twice 
the  quantity  of  H2SO4  necessary  to  neutralize  it.  Proceed  as  above.  It 
may  be  necessary  to  let  solutions  stand  a  couple  of  days  before  crystals 
appear.  Compare  crystals  as  to  appearance,  water  of  hydration,  and 
reaction  toward  litmus.  Write  equations. 

Experiment  XI. — Classification  of  Salts. 

Under  the  heads  of  Normal,  Acidic,  and  Basic,  arrange  the  following 
salts  with  reference  to  their  action  toward  litmus  paper.  A  few  cm.3 


ACIDS,  BASES  AND  SALTS.  117 

of  a  solution  of  each  will  be  sufficient.  Write  formula  for  each  salt. 
Sodium  chloride,  sodium  hydrogen  sulphate,  copper  sulphate,  aluminum 
chloride,  sodium  hydrogen  carbonate,  sodium  carbonate  and  potassium 
sulphate. 

A  salt  is  the  chief  product  of  the  interaction  of  an  acid  and  a  base,  or  (a) 
it  is  an  acid  in  which  the  hydrogen,  wholly  or  in  part,  has  been  replaced 
by  a  metal,  or  (b)  it  is  a  base  in  which  the  hydroxyl  has  been  replaced 
wholly  or  in  part  by  the  acid  radical.  Both  (a)  and  (b)  are  necessary  to 
cover  all  cases. 

This  interaction  of  acids  and  bases  is  a  species  of  reaction  known  as 
neutralization.  It  may  be  considered  as  a  process  in  which  the  removal 
from  solution  of  the  hydrogen  and  hydroxyls  is  effected,  wholly  or  in  part, 
by  arranging  the  conditions  favorable  to  their  combining  and  forming 
neutral  water.  A  more  specific  detailed  definition  may  be  offered  later. 
Neutralization  is  a  quantitative  process. 

Would  you  expect  a  definite  thermal  conduct  when  given  quantities 
of  hydrogen  and  hydroxyls  combine  to  form  molecules  of  water? 

Thermo-chemistry  of  neutralization. — 

NaOH  +  HC1  —  NaCl  +  H2O  +  137000  cal. 
KOH  +  HNO3  —  KNO3  +  H2O  +  137000  cal. 
Ca(OH)2  +  H2SO4  —  CaSO4  +  H2O  +  (2  x  137000)  cal. 

The  heat  liberated  during  the  process  of  neutralization  is  called  the 
"heat  of  neutralization." 

Acid  Salts. 
Partial, 

Basic  Salts. 
Neutralization, 

Complete,  Normal  Salts. 

An  acid  is  a  substance  containing  hydrogen  which  may  be  replaced 
wholly  or  in  part  by  a  metal. 

A  base  is  a  substance  containing  hydroxyls  which,  ordinarily,  may  be 
displaced  wholly  or  in  part  by  the  acid  radical. 

In  the  preceding  experiment  we  have  seen  that  bases  sometimes  con- 
duct themselves  like  acids. 

Pb(OH)2  +  2Na  OH  —  Na-jPb  O2  +  2H2O. 

As  we  proceed  it  will  become  evident  that  narrow  rigid  classifications 
are  not  always  possible. 

Mono-basic,  (HC1),  Mon-acid,  (NaOH), 

Basicity,  Di-basic,  (H2SO4),         Acidity,  Di-acid,  Ca(OH)2, 

Tri-basic,  (H3PO4).  Tri-acid,  Bi(OH)3. 


Il8  EXPERIMENTAL  CHEMISTRY. 

Oxacids  ending  in  ic  give  ate  salts. 

Oxacids  ending  in  ous  give  ite  salts. 

Oxacids  with  hypo  and  per  affixes  yield  corresponding  salts. 

HC1O  forms  hypochlorites, 

HC1O2  forms  chlorites, 

HC1O3  forms  chlorates, 

HC1O4  forms  perchlorates. 

A  metal  may  now  be  denned  as  a  base-forming  substance.  NH4  is 
sometimes  called  a  hypothetical  metal. 

A  non-metal  may  be  regarded  as  an  acid-forming  element.  The 
student  should  keep  clearly  in  mind  the  relations  between  hydrogen  and 
acidic  properties,  and  hydroxyls  and  basic  properties. 

PROBLEMS. 

i. — A  liter  of  a  given  solution  of  NaOH  contains  40.058  grm.  of  the 
solute;  an  acid  solution  of  unknown  strength  was  titrated  against  the  al- 
kali solution  when  it  was  found  that  50  cm.3  of  the  acid  solution  was  re- 
quired to  neutralize  25  cm.3  of  the  alkali.  What  was  the  normality 
(strength)  of  the  acid  solution? 

2. — How  many  grm.  of  H2SO4  will  be  required  to  neutralize  1500  cm.3 
of  a  2  N  solution  of  KOH?  To  neutralize  400  cm.3  of  a  5  N  solution  of 
Ba(OH)2? 

3. — If  500  grm.  of  Na-jSO^ioELjO  are  prepared  by  neutralization,  how 
many  grm.  of  acid  were  required  ?  Of  alkali  ? 

4. — If  you  were  asked  to  determine  the  strength  of  an  acid  solution, 
how  would  you  proceed,  assuming  that  you  have  access  to  a  well  equipped 
laboratory  ? 

5. — Define  Alkalimetry.     Acidimetry.     Titer.     Titration. 


CHAPTER  XIII. 

CHEMICAL  EQUILIBRIUM,  REVERSIBLE  REACTIONS,  MASS 
LAW,  DISSOCIATION. 

"A  chemical  change  in  a  given  system  is  said  to  be  complete  when 
it  proceeds  continuously  with  increase  of  'products'  and  a  correspond- 
ing decrease  of  'factors,'  until  one  or  more  of  the  factors  is  exhausted 
and  the  reaction  ceases  for  want  of  more  material.  When  increase  in 
quantity  of  products  and  decrease  in  the  quantity  of  factors  are  arrested 
before  one  or  more  of  the  factors  is  exhausted,  it  is  known  as  an  incom- 
plete chemical  change." — Ostwald. 

In  the  union  of  hydrogen  and  oxygen  at  various  temperatures,  to  form 
water,  it  will  be  recalled,  that  if  the  mixture  of  gases  (2  vols.  of  H  to  i 
vol.  of  O)  was  "sparked"  at  the  ordinary  temperature  of  the  laboratory 
the  reaction  continued  with  great  speed  until  all  the  factors  had  been 
exhausted,  i.e.,  the  reaction  ran  to  an  end.  It  will  also  be  remembered 
that  if  this  quantity  of  water  be  heated  to  a  temperature  of  2500-3000°  C., 
it  dissociates  slightly  into  H  and  O  gases.  2H2O  — *  2H2  +  O2.  Now 
it  is  further  known,  that  in  a  closed  system  composed  of  H2O  (steam) 
at  3000°  C.  and  H  and  O  gases  formed  as  the  result  of  the  decomposing 
influence  of  heat,  if  the  temperature  is  lowered  the  H  and  O  tend  to 
recombine  and  form  H2O,  that  is,  the  reaction  runs  to  the  left  as  shown 
by  the  equation  (i);  if  the  temperature  is  again  raised  then  equation  (2) 
represents  the  nature  of  the  reaction. 

(1)  2H2O  —  2H2  +  O. 

(2)  2H2O  —  2H,  +  O. 

(3)  2H20  <=>  2H2  +  O. 

It  is  evident  that  the  H  and  O  gases  are  in  equilibrium  with  water 
(steam),  and  the  reaction  is  reversible.  (Decompositions  which  are 
reversible  are  called  dissociations.)  Increase  of  temperature  tends  to 
produce  a  greater  degree  of  dissociation;  lowering  the  temperature 
tends  to  produce  a  •  greater  degree  of  association.  What  will  be  the 
effect  of  increasing  the  pressure  on  the  system  ?  Assuming  that  definite 
volumes  of  H  and  O  gases  are  in  equilibrium  with  the  steam,  apply 
pressure  or  introduce  into  the  system  volumes  of  either  gas  or  both,  then 
equation  (i)  will  represent  the  resulting  reaction;  if  the  pressure  is 
diminished  or  a  portion  of  either  gas,  or  both,  is  removed,  equation  (2) 
shows  the  direction  of  the  reaction. 

A  more  detailed  explanation  may  make  the  above  phenomena  more 
easy  of  interpretation.  Conceive  of  the  steam  at  3000°  C.  being  placed 
in  a  cylinder  which  has  a  movable  piston  and  a  stop  cock.  Keep  the 

no 


I2O  EXPERIMENTAL  CHEMISTRY. 

temperature  constant  and  diminish  the  pressure  on  the  components  by 
pulling  the  piston  well  out  toward  the  end,  the  steam  will  dissociate  in 
part  into  definite  volumes  of  H  and  O — always  the  same  quantities  under 
identical  conditions.  Inside  of  the  cylinder  there  is  a  definite  con- 
centration of  each  substance.  The  concentrations  of  the  H  and  O  gases 
are  very  small,  but  not  so  small  but  that  these  particles  meet  one  another 
every  now  and  then,  and  uniting,  form  water,  but  this  immediately 
diminishes  the  pressure  on  the  remaining  components  of  the  system 
for  three  volumes  (2  of  H  and  i  of  O)  of  hydrogen  and  oxygen  gases 
combine  to  form  only  two  volumes  of  steam.  Owing  to  this  diminished 
pressure  more  steam  dissociates  and  the  equilibrium  is  reestablished. 
Now  introduce  through  stop  cock  some  H  gas,  this  increases  the  pressure 
upon  the  system,  but  especially  does  it  increase  the  concentration  of  the 
H  particles,  with  the  result  that  they  meet  the  O  particles  more  frequently, 
and  combination  to  form  water  is  more  rapid  than  under  original  con- 
ditions. In  other  words,  the  speed  of  the  reaction  as  shown  by  (i)  is 
much  greater  than  the  speed  of  the  reaction  indicated  by  (2),  but  after 
the  reaction  (i)  has  continued  for  some  time  its  speed  will  become  less 
owing  to  diminished  concentration  of  the  hydrogen  particles.  Oxygen 
gas  introduced,  would  have  yielded  similar  results.  Any  factor  which 
alters  the  concentration  of  the  "factors"  or  "products"  will  alter  the 
general  direction  of  the  reaction.  If  either  of  the  components  of  the 
system  had  been  removed,  the  effect  upon  the  nature  of  the  ensuing 
reaction  would  have  been  just  as  marked,  only  the  reaction  would  have 
proceeded  as  indicated  by  (2). 

The  above  is  one  of  many  examples  of  concentration  effect  or  effect  oj 
mass  upon  chemical  reaction.  So  important  is  the  influence  of  "  active 
mass"  upon  chemical  reaction  that  it  has  been  stated  in  the  form  of  a 
law  by  Guldberg  and  Waage.  It  is  known  as  the  Mass  Law.* 

"  Interactions  of  substances  depend  not  only  upon  the  affinities  in- 
volved, but  also  upon  the  active  mass  oj  the  substances  in  a  unit  volume" 

2H2O  <±  2H2  +  O2. 
C2  =  C2  +  C3. 

I  2 


Ci 

K  is  sometimes  called  the  affinity  constant.     It  really  represents  the 
ratio  of  the  affinities  urging  the  opposed  actions.     The  speed  of  the 

reaction  represented  by  (i)  may  be  indicated  by  Speed!  =  Kj  C  i  and 
the  opposed  reaction  by  Speed2  =  K2C^.C3  where  Kx    represents  the 

*  Professor  T.  W.  Richards  prefers  to  call  it,  "  The  Law  of  Concentration  Effect. 


CHEMICAL    EQUILIBRIUM,    REVERSIBLE    REACTIONS,    ETC.  121 

tendency  for  water  to  dissociate,  and  K2  the  tendency  for  H  and  O  to 
combine. 

K,         C22.C3  C22.C3 

K.C'-K.C'.C,;-  -;K  = 

K2        C*  C* 

Experiment  I. — Reactions  which  are  Approximately  Complete. 

(a)  Formation  of  an  inactive  and  practically  insoluble  product.     Pour 
3  cm. 3  or  4  cm.s  of  silver  nitrate  (AgNO3)  solution  into  a  test  tube, 
(use  side-shelf  reagent);  add  a  few  drops  of  a  dilute  solution  of  sodium 
chloride  (NaCl);  let  the  precipitate  curd  and  settle  to  the  bottom;  observe 
whether  an  increase  in  the  quantity  of  the  precipitant  will  cause  additional 
precipitation.     Continue  to  add  the  precipitant  until  subsequent  addi- 
tions do  not  effect  further  precipitation  of  silver  chloride  (AgCl).     Equa- 
tion?    Are  the  "products"  of  the  reaction,  active  or  inactive?     Explain. 
Is  it  possible  that  under  certain  conditions  the  "products"  may  react 
to  produce  the  "factors?"     Are  all  substances  soluble  in  degree?     Is 
the  precipitate  (AgCl)  in  equilibrium  with  the  small  quantity  which  is 
dissolved  ? 

(b)  Formation   of   a  product  (gaseous)  which  is   removed.     Prepare 
a  dilute  solution  of  sodium  carbonate  (Na^CC^);  test  with  litmus  solution; 
add  a  few  drops  of  dilute  sulphuric  acid  (H2SO4)  at  a  time  and  note  effect 
upon  the  speed  of  the  reaction;  continue  to  add  acid  until  the  effervescence 
of  CO2  has  ceased,   and  blue  litmus  begins  to  turn  red.     Equation? 
When  one  of  the  products  escape  from  the  system,  is  it  likely  that  the 
"products"  will  react  to  produce  the  so-called  "factors?"     Is  there  any 
likelihood  of  the  establishment  of  an  equilibrium  between  the  "factors" 
and  the  "products?"    If  not,  is  the  reaction  reversible?     Explain. 

Experiment  II. — Reactions  which  are  Incomplete.     Equilibrium. 

(a)  Reversible  reaction.  To  5  cm. 3  of  a  magnesium  sulphate  or 
chloride  solution  (side-shelf),  add  a  few  drops  of  NH4OH  at  a  time: 
observe  the  partial  precipitation  of  magnesium  hydroxide  (Mg(OH)2); 
now  add  an  excess  of  NH4OH  and  note  that  subsequent  additions  of 
NH4OH  do  not  effect  further  precipitation. 

(i)  MgCl2  +  2NH4OH  <=*_M_g(OH)_2  +  2NH4C1. 

Now  add  an  excess  of  ammonium  chloride,  NH4C1 — one  of  the  "pro- 
ducts," and  observe  that  the  presence  of  a  large  concentration  of  this 
"active  mass"  causes  reversal  of  reaction  as  indicated  by  (i). 

Mg(OH)2  +  NH4C1  (excess)  -»  MgCl2.2NH4Cl  +  NH4OH 
MgCl2  +  NH4OH  +  NH4C1  (excess)  ->  MgCl2.2NH4Cl  +  NH4OH. 

Inasmuch  as  Mg(OH)2  is  soluble  in  NH4C1  which  is  formed  simul- 
taneously, would  you  say  that  there  is  a  complete  precipitation  in  the 


122  EXPERIMENTAL  CHEMISTRY. 

first  reaction?  Do  the  "products''  tend  to  react  and  reproduce  the 
" factors?"  When  the  speeds  of  these  opposing  reactions  are  the  same, 
a  condition  of  equilibrium  results.  Was  the  speed  of  the  reverse  reaction 
increased  by  increasing  the  concentration  of  one  of  the  "  active  products  "  ? 
If  the  NH4C1  is  removed  as  rapidly  as  it  is  formed  in  first  reaction,  will 
the  reaction  be  complete  ?  Will  there  ensue  a  condition  of  equilibrium  ? 
Will  the  reaction  be  reversible? 

(b)  An  incomplete  reaction  completed  by  rendering  one  of  the  active 
products  inactive  by  virtue  of  its  insolubility  in  another  medium.     To 
15  cm.3  of  a  dilute  solution  of  calcium  chloride,  CaCl2,  add  a  few  cm.3 
of  a  concentrated  solution  of  potassium  sulphate,  K2SO4;  observe  the 
slight  and  partial  precipitation  of  calcium  sulphate,  CaSO4. 

CaCl2  +  K2SO44=±  CaSO4  +  2K  Cl. 

Filter,  and  add  5  cm.3  of  ethyl  alcohol,  C2H5OH,  to  5  cm.3  of  the 
filtrate.  Explain  the  formation  of  the  heavy  precipitate  of  CaSO4. 

(c)  A  reaction  involving  equilibrium,  in  which  each  of  the  opposing 
reactions  are  practically  completed  by  varying  the  concentration  of  the 
"active  masses." 

Add  a  few  cm.3  of  an  oxalic  acid,  H2C2O4,  solution  to  5  cm.3  of  a  dilute 
CaCl2  solution;  the  precipitation  of  the  calcium  oxalate,  CaC2O4,  is 
incomplete  owing  to  the  reaction  of  HC1  which  is  formed  simultaneously. 

CaCl2  +  H2C2O4<=>CaC2O4  +  2HC1. 

Decant  half  of  the  solution  upon  a  filter,  and  collect  the  filtrate.  To 
the  portion  remaining  in  the  test  tube  add  HC1  until  precipitate  dis- 
solves. Write  equation.  To  the  filtrate  add  two  or  three  cm.3  of  a 
sodium  acetate,  NaC2H3O2,  solution.  Explain  formation  of  precipitate 
of  CaC2O4.  Write  equations.  Compare  the  reactions  as  represented 
by  equations.  Are  the  reactions  as  represented  by  your  equations, 
opposite  in  nature?  If  these  two  opposing  reactions  should  occur  in 
one  system,  with  identical  speeds,  would  a  condition  of  equilibrium 
results  ? 

Berthollet's  Laws  are  in  substance  as  follows: 

"  //  the  '  products '  are  not  active,  or  are  active  but  are  removed,  then  the 
reaction  is  complete  and  runs  to  an  end." 

"  When  the  products  of  a  reaction  are  chemically  active  within  a  system 
and  are  not  removed,  the  reaction  is  reversible  and  incomplete,  and  results 
ultimately  in  balanced  action  and  chemical  equilibrium." 

Chemists  are  now  inclined  to  regard  all  chemical  changes  as  being 
reversible,  i.e.,  reactions  do  not  run  to  completion. 

A  thorough  understanding  of  Le  Chatelier's  Theorem  will  enable 
the  student  to  understand  much  in  regard  to  chemical  equilibrium  which 
would  otherwise  be  vague.  (See  Energetics  of  Chemistry.) 

Before  leaving  this  subject   the  student  is  warned  against  looking 


CHEMICAL    EQUILIBRIUM,    REVERSIBLE    REACTIONS,  ETC.  123 

upon  a  condition  of  chemical  equilibrium  as  being  due  to  a  cessation  or 
suspension  of  chemical  action.  On  the  contrary  such  a  state  of  balance 
is  due  to  the  fact  that  the  speeds  of  the  opposing  reactions  are  equal. 
Again,  a  reversible  action  must  not  be  considered  as  one  which  runs  to 
completion,  when  an  opposite  reaction  is  then  initiated  and  runs  back- 
ward. Both  actions  are  started  at  practically  the  same  time;  the  one 
gradually  increases  in  speed  and  the  other  gradually  diminishes  in  speed 
until  finally  their  speeds  are  identical,  when  a  condition  of  equilibrium 
results. 

Experiment  III. — Dissociation  of  a  Gas. 

Place  a  small  quantity  of  NH4C1  in  a  dry  test  tube;  heat  tube  until 
white  fumes  are  evolved  and  sublime  up  the  tube.  Now  hold  a  piece 
of  moistened  red  litmus  paper  in  the  mouth  of  the  tube  and  it  will  be 
turned  blue  showing  that  a  portion  of  the  ammonia  is  escaping;  discon- 
tinue heating  tube  for  a  moment  and  the  blue  litmus  paper  will  be  turned 
red  owing  to  the  hydrogen  chloride  which  is  now  escaping.  Heat  tube 
strongly,  and  hold  litmus  paper  in  the  heavy  white  fumes  which  appear 
a  few  inches  above  the  mouth  of  the  tube.  Results  ?  This  is  a  case  of 
dissociation.  Is  temperature  a  factor  in  determining  which  way  the 
reaction  shall  go  ?  Explain.  Would  you  say  the  heat  oj  dissociation  is 
positive  or  negative  ? 

NH4C1  <=>  NH3  +  HCl 

Many  gases  tend  to  dissociate  when  placed  under  favorable  energy 
conditions.  At  a  temperature  of 

2HI  <=±  H2  +  I2 
N2O4  <=»  NO2  +  NO2 
I2^I  +  I 
NH4C1  <±  NH3  +  HCl. 

480°  C.  about  20  per  cent,  of  the  original  quantity  of  HI  is  in  the  dis- 
sociated condition. 


CHAPTER  XIV. 
NOTE  ON  THE  MODERN  THEORY  OF  SOLUTION. 

The  subject  of  "solutions"  was  introduced  in  a  preceding  chapter, 
but  little  was  said  about  the  condition  of  the  solute  when  in  solution. 
Attention  was  called  to  the  fact  that  a  substance  in  solution  conducts 
itself  very  much  like  a  gas,  and  when  the  dissolved  portion  is  in  equilib- 
rium with  an  undissolved  portion,  the  solution  is  said  to  be  saturated. 
However,  the  mechanism  of  a  solution  was  not  considered  in  detail. 

The  purpose  of  these  notes  and  the  following  experiments  is  to  present 
a  comprehensive  view  of  the  modern  theory  of  solution. 

The  theory  has  grown  from  the  wrecks  of  other  theories.  Glimpses  of 
and  approximations  to  the  present  theory  logically  preceded  it.  There- 
fore, it  is  somewhat  difficult  to  say  just  where  it  had  its  beginnings;  how- 
ever, it  is  certainly  the  result  of  bringing  together,  interpreting  and  corre- 
lating a  mass  of  apparently  unrelated  experimental  data. 

Without  detracting  one  iota  from  the  man  who  elaborated  the  theory 
in  practically  its  present  form,  we  must  acknowledge  that  we  are  indebted 
to  other  men  as  well  as  to  Arrhenius  for  the  "Ionic  Hypothesis"  or  The 
Electrolytic  Dissociation  Theory. 

Although  the  theory  in  its  present  form  has  been  unable  to  meet  in 
an  altogether  satisfactory  manner  all  of  the  requirements  made  of  it, 
yet  it  is  quite  possible  when  various  relationships  are  more  perfectly 
understood  that  it  will  meet  all  demands.  It  is  so  vastly  superior  in  its 
ability  to  interpret  an  array  of  phenomena  which  would  otherwise  be 
inexplicable  that  many  noted  investigators  have  called  the  theory  "a 
corner-stone  of  physical  chemistry." 

On  investigating  the  labors  of  the  chemists  and  physicists  who  pre- 
ceded and  were  contemporaneous  with  Arrhenius,  it  appears  as  though 
it  would  be  perfectly  justifiable  to  affirm  that  the  theory  as  elaborated 
was  or  could  have  been  reached  along  practically  two  -independ- 
ent lines  of  thought  and  investigation;  namely,  "electro-chemical" 
and  "osmotic  and  vapor  pressure  effects."  That  as  it  may  or  may  not 
be,  the  literature  reveals  Arrhenius,  himself  working  along  electro-chemi- 
cal lines,  as  standing  at  the  apex  of  two  converging  lines  of  investigation; 
in  fact,  responsible  for  causing  the  lines  to  meet  when  they  did.  It  was 
Arrhenius  who  elaborated  and  enunciated  the  theory  which  correlated 
a  great  volume  of  isolated  data.  Other  men  had  been  unable  to  do  so. 
It  was  Ostwald  who  then  took  up  the  theory,  applied  it,  and  finally  be- 
came its  most  effective  and  influential  exponent.  The  theory  now  meets 
with  an  almost  universal  acceptance. 

To  use  the  historical  method  of  approach  to  our  subject — that  is,  to 

124 


NOTE    ON    THE    MODERN    THEORY    OF    SOLUTION.  125 

move  historically  along  the  two  lines  above  enumerated — would  be  per- 
haps the  most  satisfactory,  but  this  work  is  neither  so  voluminous  nor 
pretentious  as  to  attempt  more  than  to  merely  mention  briefly  the  names 
of  the  men  who  have  contributed  to  the  development  of  the  theory. 

Electro-Chemical. — We  shall  not  comment  on  the  labors  of  Gilbert, 
Dufay,  Becaria,  Priestly,  Cavendish,  Franklin,  Wilke,  Galvani,  Volta  and 
Ritter,  J.  W.,  further -than  to  say  that  Volta  originally  separated  all  con- 
ductors of  electricity  into  two  classes — a  first  class,  comprising  the  metals, 
carbon  and  other  good  conducting  substances  found  in  nature,  such  as 
metallic  sulphides;  the  second  class  embraced  such  solutions  as  are  con- 
ductors of  electricity.  Until  very  recently  we  have  spoken  of  the  first  class 
of  conductors  as  those  in  which  when  the  current  is  flowing  there  is  no 
"simultaneous  motion  of  the  ponderable  matter,"  while  in  the  case  of  the 
second  class  there  is  supposed  to  be  a  "  corresponding  motion  of  ponder- 
able matter."  That  the  flow  of  electricity  through  conductors  of  the  first 
class  is  convectional,  as  well  as  in  the  case  of  conductors  of  the  second 
class,  is  a  recent  theory  that  is  supported  by  many. 

About  the  year  1800  Nicholson  and  Carlisle  found  that  when  an  electric 
current  was  passed  through  two  wires  whose  ends  dipped  into  water,  that 
hydrogen  and  oxygen  gases  were  liberated,  and  further,  the  liquid  around 
the  pole  at  which  the  hydrogen  was  evolved  was  alkaline,  while  it  became 
acid  around  the  other  pole. 

In  1802  Ermann  secured  results  very  similar  to  those  of  Nicholson  and 
Carlisle.  Men  attempted  to  explain  the  acid  and  alkaline  reactions  by 
asserting  that  it  was  due  to  the  effect  of  electricity  upon  water. 

Sir  Humphry  Davy  (1.778-1829)  investigated  the  above  phenomena 
and  demonstrated  that  pure  water  is  separated  into  hydrogen  and  oxygen 
gases,  and  that  the  acid  and  alkaline  properties  were  due  to  impurities  in 
the  water.  Davy  performed  the  experiment  known  to  us  as  the  "electroly- 
sis of  sodium  sulphate,"  the  solution  having  previously  been  colored  blue 
with  a  litmus  solution.  A  satisfactory  explanation  of  the  phenomena  was 
not  offered  until  many  years  later.  Davy  advanced  that  which  may  be 
called  the  first  "electro-chemical  theory."  It  was  based  upon  Dalton's 
atomic  hypothesis.  In  brief,  his  theory  assumed  that  the  atoms  of  the 
elements  when  they  came  in  contact  with  one  another  took  to  themselves 
opposite  charges  of  electricity;  that  if  the  charges  were  sufficiently  strong 
the  two  elements  combined,  forming  a  chemical  compound.  Also,  if  a 
new  and  different  atom  came  in  contact  with  the  compound  and  could 
assume  a  sufficiently  strong  charge  of  electricity,  it  would  attract  the  oppo- 
sitely charged  atom  of  the  compound,  causing  decomposition.  The 
theory  did  not  meet  with  a  ready  or  even  an  ultimate  acceptance. 

Berzelius  (1779-1848)  now  advanced  his  "dualistic  theory"  which  in- 
volved the  "bipolarity  of  atoms."  Affinity  depended  upon  the  size 
and  sign  of  the  dominant  electrical  charge  on  the  atom.  Every  com- 
pound was  conceived  of  as  being  composed  of  two  parts,  oppositely 
charged — the  parts  themselves  might  be  composed  of  two  parts.  For 
example,  the  salt,  Zn  SO4,  should  be  written  according  to  Berzelius,  ZnO.- 


126  EXPERIMENTAL  CHEMISTRY. 

SO3,  to  show  its  real  structure  as  being  composed  of  a  basic  oxide  and  an 
acid  anhydride.  (See  Le  Blanc's  Electro-Chemistry.) 

Grotthus  (1805)  announced  the  first  complete  theory  of  electrolysis, 
known  as  the  "exchange  of  partner  theory."  Clausius  showed  theory 
to  be  a  violation  of  the  "second  law  of  energetics."  (Lehfeldt's  Electro- 
Chemistry  contains  a  concise  statement  of  the  theory.) 

Faraday  (1794-1868)  found  that  the  quantity  -of  electricity  passing 
through  a  circuit  and  the  chemical  and  magnetic  effects  produced  were 
proportional  to  one  another.  By  a  series  of  careful  measurements  Fara- 
day, in  1833,  established  the  first  of  two  laws  known  as  Faraday's  Laws.  In 
substance  the  First  Law  is:  the  amounts  of  substances  which  separate  at  the 
electrodes  during  electrolysis  are  strictly  proportional  to  the  quantity  of 
the  electric  current  which  passes  through  the  electrolyte.  The  Second  law 
says:  the  mass  of  any  substance  liberated  at  the  electrodes  by  a  given 
quantity  oj  electricity  is  directly  proportional  to  the  " chemical  equivalent" 
or  combining  weight.  Davy  believed  that  the  current  was  carried  through 
solutions  by  little  particles  called  "ions."  He  also  injected  into  our  nom- 
enclature such  terms,  as  electrolysis  (the  process  of  decomposing  sub- 
stances in  solution  by  an  electric  current),  electrolyte  (the  substance  which 
carries  the  current),  anode  (the  positive  electrode),  cathode  (the  negative 
electrode),  ions  (the  particles  into  which  the  substance  divides  and  act 
as  the  real  carriers  of  the  current),  anions  (the  ions  which  gather  at  the 
anode),  and  the  cations  (ions  which  gather  at  the  cathode). 

Note. — The  student  is  reminded  that  although  we  shall  retain  these 
terms,  Faraday  did  not  use  them  in  the  sense  in  which  they  are  used  at  the 
present  time. 

At  that  time  there  was  considerable  discussion  as  to  what  constituted 
the  anions  and  the  cations.  Berzelius  said  in  the  case  of  sodium  sulphate, 
which  he  wrote  NaO  and  SO3,  that  NaO  and  SO3  constituted  respectively 
the  cations  and  anions — that  these  reacted  upon  the  water  and  produced  an 
alkali,  NaO  +  H2O->  Na(OH)2,  and  an  acid,  SO3  +  H2O  —  H2SO4, 
but  it  was  known  that  hydrogen  and  oxygen  were  liberated  at  the  same 
time. 

Daniell  disproved  Berzelius's  explanation  by  experimentally  showing 
that  the  hydrogen  and  oxygen  liberated  were  chemically  equivalent  to  the 
alkali  and  acid  formed.  This  would  require,  so  to  speak,  "  double  electrical 
action"  which  would  be  contrary  to  Faraday's  laws.  Daniell  affirmed 
that  Na  was  the  cation  and  SO4  the  anion;  that  these  ions  were  set  free  at 
the  poles,  and  then  reacted  upon  the  water  liberating  hydrogen  and  oxy- 
gen. "  The  salt  alone  must  have  conducted  the  electricity  in  the  solution," 
for  if  the  water  had  conducted  a  part  of  the  current  of  electricity  there 
would  not  have  existed  the  interequivalence  of  acid  and  alkali,  and 
hydrogen  and  oxygen. 

It  became  known  later  that  both  Ohm's  law  and  Joule's  law  hold  for 
conductors  of  the  second  class  as  well  as  for  the  first  class,  therefore  it 
was  obvious  that  none  of  the  current  spent  in  traversing  an  electrolyte 


NOTE    ON   THE   MODERN    THEORY    OF    SOLUTION.  127 

was  employed  in  doing  chemical  work  in  splitting  up  the  electrolyte  into 
ions.  Within  the  electrolyte  during  electrolysis  the  current  merely  exerts 
a  directive  force  on  the  ions;  that  is,  so  to  speak,  " sorts"  them. 

The  theory  of  Grotthus,  as  mentioned  previously,  was  found  to  be  in 
conflict  with  the  " second  law  of  energetics"  by  Clausius,  who  advanced 
a  theory  of  his  own,  in  which  he  set  forth  the  idea  that  the  positive  and 
negative  parts  of  a  molecule  in  a  solution  are  frequently  in  such  a  rapid 
state  of  vibration  that  for  a  few  moments  at  a  time  the  parts  may  be  said 
to  be  independent  of  one  another,  so  that  if  a  current  were  passing  at  the 
time  the  parts  might  be  so  favorably  arranged  as  to  follow  the  directive 
force  of  the  current  and  be  deposited  on  electrodes.  In  other  words, 
he  assumed  a  constant  interchange  of  the  negative  and  positive  parts  of 
the  molecule,  and  at  any  instant  there  was  a  small  number  of  these  parts 
of  the  molecules  momentarily  free,  and  that  the  current  was  carried  by 
these  molecule-parts.  This  theory  was  commonly  accepted  and  is  the 
one  which  preceded  our  present  theory. 

Hittorf  began  his  work  about  this  time,  on  what  he  called  the  "  migration 
of  the  ions."  He  found  during  the  electrolysis  of  a  CuSO4  solution  be- 
tween electrodes  of  copper,  that  the  solution  became  very  much  more  con- 
centrated around  the  anode  then  around  the  cathode,  yet  the  quantity  of 
copper  deposited  at  the  cathode  was  greater  than  the  amount  lost  by 
the  cathode  chamber.  The  conclusion  was  not  difficult  that  the  ions  of 
Cu  and  SO4  had  different  migration  velocities.  From  the  speed  ratios 
Hittorf  calculated  the  "transport  numbers." 

Kohlrausch,  as  the  result  of  his  work  on  the  conductivity  of  dilute  solu- 
tions, established  a  very  simple  relation  between  the  "  transport  numbers  " 
and  the  molecular  conductivity.  Kohlrausch's  work  also  emphasized  the 
fact  that  inorganic  substances,  such  as  acids,  bases  and  salts,  were  good 
conductors  (electrolytes),  and  organic  substances  were  either  non-conduc- 
tors or  very  poor  conductors. 

As  the  result  of  the  labors  of  Hittorf  and  Kohlrausch,  Svante  Arrhenius 
(1887),  a  Swedish  chemist,  was  enabled  to  reach  certain  conclusions  re- 
garding the  theory  of  "free  ions."  Arrhenius  while  working  on  the  con- 
ductivity of  solutions  recognized  two  kinds  of  "molecules" — one  active 
and  the  other  inactive — as  he  named  them  because  of  the  belief  that  only 
the  "active"  molecule  was  instrumental  in  causing  conductivity.  He 
stated  further  his  belief  that  the  "inactive"  changed  into  "active" 
molecules  with  sufficient  dilution.  He  was  not  able  to  offer  a  convincing 
proof  until  he  was  placed  in  receipt  of  certain  data  from  the  other  line 
of  investigation  referred  to  before,  namely,  "  osmotic  and  vapor  pressure  " 
effects. 

It  had  not  been  so  very  long  before  this  time  that  investigators  had  ac- 
cepted with  considerable  hesitancy  "the  dissociation  of  molecules"  as 
the  explanation  of  the  deviations  of  many  vapor  densities  from  that  antici- 
pated by  theory.  Planck  at  this  time  as  the  result  of  purely  thermody- 
namical  considerations  said  that  "it  would  be  quite  natural  to  expect  a 
similar  dissociation  when  substances  were  placed  in  solution." 


128  EXPERIMENTAL  CHEMISTRY. 

Osmotic  Pressure. — Without  attempting  to  make  a  complete  review  of 
all  the  work  that  had  been  done  along  this  line  up  to  the  time  of  Arrhen- 
ius,  it  will  be  dismissed  with  the  statement  that  the  question  of  solution 
had  been  holding  the  attention  of  eminent  investigators  for  a  period  of 
ten  to  fifteen  years  preceding  the  time  at  which  Arrhenius  announced  his 
theory.  The  fundamental  idea  which  stimulated  men  to  these  investi- 
gations was  that  when  a  substance  is  in  the  dissolved  condition  it  behaves 
in  a  manner  very  similar  to  a  gas.  Although  the  idea  was  not  new,  yet  it 
secured  the  attention  of  Van't  Hoff  (now  in  the  University  of  Berlin), 
who  became  firmly  convinced  of  the  fact  that  a  "  substance  in  solution 
behaves  like  a  gas  and,  like  it,  exerts  a  pressure"  but  he  was  unable  to 
devise  a  method  for  the  measurement  of  the  pressure. 

We  are  told  that  Van't  Hoff  was  one  day  walking  with  De  Vries,  the 
zoologist,  and  remarked  about  the  pressure  which  forces  upward  the  sap 
in  a  tree,  whereupon  De  Vries  replied  that  Pfeffer,  the  botanist,  had  in- 
vestigated and  also  measured  the  force.  This  was  the  very  data  for 
which  Van't  Hoff  had  been  searching. 

Pfeffer  had  observed  some  years  previously  that  when  a  plant  cell  was 
placed  in  very  dilute  solutions  it  bursted,  and  when  placed  in  concen- 
trated solutions,  the  cell  shriveled.  At  the  time  Pfeffer  was  unable  to 
make  any  satisfactory  measurements,  but  later,  due  to  the  labors  of  Traube, 
he  was  able  to  deposit  semipermeable  membranes  of  copper-ferrocyanide 
in  the  walls  of  porous  porcelain  cups  which  he  connected  with  a  mano- 
meter; the  cup  was  then  filled  with  a  sugar  solution  and  immersed  in 
pure  water.  As  the  result  of  Pfeffer's  experiments,  although  he  was  una- 
ware of  their  far-reaching  influence,  it  was  found  that  the  pressure  (osmo- 
tic) was  directly  proportional  to  the  concentration  and  the  absolute  tem- 
perature of  the  solution.  Van't  Hoff  was  quick  to  correlate  osmotic 
pressure  and  gas  pressure.  He  found,  in  other  words,  that  the  gas  laws 
hold  approximately  for  osmotic  pressure.  Further,  in  substance,  that 
equal  fractions  of  the  molecular-gram  weights  of  organic  substances  dis- 
solved in  the  same  volume  of  water  produce  the  same  osmotic  pressure. 
To  state  his  conclusions  from  another  point  of  view  and  include  Avo- 
gadro's  rule,  "At  the  same  osmotic  pressure  and  temperature  equal 
volumes  of  all  solutions  contain  the  same  number  of  molecules,  and,  in 
fact,  that  number  which  under  the  same  pressure  and  at  the  same  temperature 
exists  in  the  same  volume  of  gas."  As  Van't  Hoff  proceeded  with  his  work 
he  found  that  strong  inorganic  acids  and  bases  and  salts  did  not  act  nor- 
mally, i.e., the  osmotic  pressures  as  yielded  by  solutions  of  these  substances 
were  sometimes  more  than  double  that  which  he  anticipated;  for  example, 
when  a  given  weight  of  hydrogen  chloride  was  dissolved  in  water  it  gave 
an  osmotic  pressure  equal  to  almost  twice  that  which  it  should  yield  ac- 
cording to  the  gas  law  (PV  =  RT).  He  was  unable  to  account  for  this. 

Freezing  Point  and  Vapor  Pressure  of  Solvents. — About  this  time 
Raoult  made  two  important  generalizations  in  regard  to  dilute  solutions 
which  came  as  the  result  of  the  labors  of  himself,  Bab,  Wullner,  Ostwald, 
Blagden,  Rudolph  and  others.  Our  knowledge  of  the  depression  of  the 


NOTE    ON    THE   MODERN    THEORY    OF    SOLUTION.  I2Q 

freezing  point  and  elevation  of  the  boiling  points  of  solvents  by  dissolved 
substances,  was  uncertain  and  fragmentary  before  Raoult  initiated  his 
investigations.  (See  chapter  on  "Solutions.") 

General  Law  of  the  Vapor  Pressure  of  Solvents. — i  molecule  oj  a  non- 
saline,  non-volatile  substance,  dissolved  in  100  molecules  of  any  volatile 
liquid,  lowers  the  vapor  pressure  oj  this  liquid  by  a  nearly  constant  fraction 
oj  its  value — approximately  .0105. 

The  experience  of  Van't  Hoff  relative  to  osmotic  pressures  was  the 
experience  of  Raoult  in  regard  to  deviations  from  his  law  concerning  the 
freezing  point,  namely,  that  organic  compounds  gave  normal  and  com- 
paratively uniform  results,  while  strong  inorganic  acids  and  bases  and 
salts  gave  a  larger  depression  than  was  anticipated.  Raoult  was  like- 
wise unable  to  explain  it. 

In  the  year  1887  Van't  Hoff  contributed  an  article  to  the  Zeitschrift  fiir 
physikalische  Chemie  on  "  The  Role  of  Osmotic  Pressure  in  the  Analogy 
Between  Solutions  and  Gases."  The  article  was  in  support  of  his  be- 
lief that  the  gas  law  was  applicable  to  dilute  solutions  of  substances. 
It  called  attention  to  the  work  of  Rauolt  and  emphasized  the  fact  that 
the  use  of  strong  acids  and  bases  caused  deviations  from  both  his  own 
and  Raoult's  laws.  Near  the  latter  end  of  the  article  Van't  Hoff  says, 
"  *  *  *  Arrhenius  pointed  out  to  me,  by  letter,  the  probability  that  salts 
and  analogous  substances  when  in  solution  break  down  into  ions.  As  a 
matter  of  fact,  as  far  as  investigation  has  been  carried,  the  solutions  which 
obey  the  law  of  Avogadro  are  non-conductors,  which  indicates  that  they 
are  not  broken  down  into  ions;  and  a  further  experimental  examination 
of  other  solutions  is  possible,  since,  from  the  assumption  made  by  Arrhen- 
ius, the  deviation  from  Avogadro's  law  can  be  calculated  from  the  con- 
ductivity." 

As  soon  as  the  above  article  appeared,  Arrhenius  was  able,  by  comparing 
the  effects  of  electrolytes  and  non-electrolytes  in  depression  of  the  freezing 
point  of  water  with  their  respective  electrical  conductivities,  to  produce 
convincing  proof  of  his  original  assumption  of  the  electrolytic  disso- 
ciation of  certain  salts  in  aqueous  solution.  The  organic  compounds  of 
Van't  Hoff  and  Raoult  were  the  non-electrolytes  of  Kohlrausch  and 
Arrhenius. 

Arrhenius  set  forth  his  proof  and  theory  in  an  article  entitled,  "  Uber 
die  Dissociation  der  im  Wasser  gelosten  Stoffe."  He  called  attention  to 
the  fact  that  the  compounds,  salts,  strong  acids  and  strong  bases  which 
create  abnormally  high  osmotic  pressures  and  abnormal  depressions  of 
the  freezing  point  are  all  electrolytes — i.  e.,  when  dissolved  in  water 
their  solutions  become  good  conductors  of  electricity;  on  the  other  hand, 
the  substances  which  give  normal  osmotic  pressures  and  normal  depres- 
sions of  the  freezing  points  are  either  non-electrolytes  or  conduct  the 
current  very  poorly.  A  normal  solution  of  cane  sugar,  which  is  a  non- 
electrolyte,  freezes  at  — 1.87°  C.;  a  normal  solution  of  sodium  chloride 
which  is  a  good  electrolyte  freezes  at  — 3.46°  C.  A  very  similar  relation 
holds  in  regard  to  the  osmotic  pressures  of  the  two  solutions.  It  is  gener- 
9 


130  EXPERIMENTAL  CHEMISTRY. 

ally  accepted  that  the  pressure  of  a  gas  or  the  osmotic  pressure  or  the  de- 
pression of  the  freezing  point  is  directly  proportional  to  the  number  of 
molecules  or  ultimate  particles  present  in  the  gas  or  solution.  Arrhenius 
concluded,  therefore,  that  if  electrolytes  produced  abnormal  alterations  in 
these  particular  properties,  that  the  molecules  of  the  substance  in  solution 
split  up  (dissociated)  into  a  larger  number  of  particles  which,  so  far  as 
above  phenomena  are  concerned,  acted  like  molecules.  Each  of  these 
particles  would  produce  its  own  effect  on  osmotic  pressure  or  depression  of 
the  freezing  point  and  that  inasmuch  as  those  substances  which  behaved 
normally  with  reference  to  osmotic  pressure,  etc.,  were  non-electrolytes, 
these  smaller  particles  were  the  real  cause  of  the  conductivity  of  a  solution. 
Arrhenius  gave  the  general  name  ions  to  these  independent  particles  into 
which  molecules  were  assumed  to  dissociate.  The  terminology  of  the 
theory  is  very  similar  to  that  suggested  by  Faraday;  in  fact,  many  parts 
are  identical. 

Theory. — In  briefest  outline,  his  theory  is  as  follows:  Whenever  an  elec- 
trolyte, referring  to  the  solute,  is  dissolved  it  almost  wholly  or  partly  disso- 
ciates into  ions.  The  extent  of  the  dissociation  depends  upon  nature  of  so- 
ute,  concentration,  solvent,  temperature  and  pressure.  These  ions  many  be 
composed  of  one  or  more  atoms,  but  are  altogether  different  in  nature  from 
the  elementary  substances  of  which  they  are  composed.  These  ions 
are  electrically  charged,  whereas  the  elementary  substances  are  neutral, 
therefore  their  energy  content  is  different,  and  logically  we  expect  a  differ- 
ence in  properties.  An  electrolyte  dissociates  into  two  ions — one  (cation), 
bearing  a  positive,  and  the  other  (anion),  a  negative  charge  of  electricity. 
Frequently  one  of  the  ions  further  dissociates  or  it  may  associate  with 
another  ion  forming  a  complex  ion.  It  should  be  remembered  that  the 
solution  as  a  whole  is  electrically  neutral  and  therefore,  there  must  be  as 
many  electrical  charges  of  one  kind  as  there  are  of  another.  Irrespective 
of  the  masses  of  the  ions,  the  number  of  charges  carried  by  each  ion  is 
equal  to  the  valency  of  the  atom  or  atomic  group  which  constitutes  the 
ion.  As  a  matter  of  fact,  on  the  electrical  theory,  valency  really  amounts 
to  nothing  more  or  less  than  the  "number  of  positive  or  negative  unit 
charges  associated  with  the  chemical  atom."  An  example  may  make  the 
theory  more  easily  understood.  When  hydrogen  chloride  is  dissolved  in 
water  it  immediately  dissociates  into  ions  of  hydrogen  (hydrion)  and  chlor- 
ine (chloridion),  respectively  known  as  cations  and  anions.  The  ions  are 
wholly  independent  of  one  another.  Each  ion  retains  its  identity,  i.e., 
properties  of  reaction,  velocity,  etc.,  regardless  of  its  birth.  Dilution  in- 
creases the  extent  of  the  dissociation.  If  the  solution  is  examined,  there 
will  be  no  evidence  of  any  free  atoms  of  either  hydrogen  or  chlorine  gas. 
If  these  gases  were  in  the  free  state  they  would  be  very  easy  to  detect  as 
hydrogen  is  very  slightly  soluble  and  chlorine  is  a  yellow  colored  gas. 
The  fact  that  no  electrical  energy  is  consumed  in  dissociating  the  mole- 
cule within  the  liquid,  and  yet  when  a  current  is  passed  through  the  solu- 
tion the  hydrogen  ions  (cations)  are  directed  toward,  and  will  appear,  at 
the  negative  electrode  (cathode),  while  the  chlorine  ion  (anion)  is  attracted 


NOTE    ON    THE    MODERN    THEORY    OF    SOLUTION.  131 

to  the  positive  electrode  (anode),  is  strong  evidence  that  the  electrolyte 
immediately  dissociates  on  being  placed  in  solution  and  that  the  ions 
are  the  actual  carriers  of  the  current.  When  the  ions  reach  their  respective 
electrodes  their  charges  of  electricity  are  unloaded  and  they  are  trans- 
formed into  the  elementary  substances;  they  are  then  deposited  upon  the 
electrode,  escape  as  a  gas  or  react  chemically  upon  the  water.  Since 
hydrogen  has  a  valency  of  one  it  is  said  to  have  one  electrical  charge 
and  that  positive;  chlorine 

HC1,  aq.  <=±  H  +  Cl'  NaOH,  aq.  +±  Na  +  OH' 

ions  have  one  negative  charge. 

H2SO4,  aq.  <±  H   +  HSO'4        NaCl,  aq.<=±  Na  +  Cl' 

lt  + 
S0"4  +  H  . 

It  is  evident  that  in  an  electrolytic  solution  the  solute  is  partly  in  the  molec- 
ular condition  and  partly  in  the  ionic  state.  Water  is  not  a  good  con- 
ductor of  electricity,  yet  it  is  very  slightly  dissociated. 

(99-95%)H20^±H  +  +  OH'(.o4i%) 

The  remark  is  frequently  made,  however,  that  if  it  is  slightly  acidified  with 
sulphuric  acid  it  becomes  a  good  conductor.  Interpreted  in  terms  of 
the  "electrolytic  dissociation"  theory,  this  means  merely  that  the  acid  is 
dissociated  into  hydrogen  and  the  acid  radical  ions  which  carry  the 
current. 

The  student  should  not  confuse  the  ideas  of  solubility  and  electrolytic 
dissociation.  Great  solubility  does  not  imply  great  dissociation. 

For  purposes  of  convenience  all  substances  have  been  divided  into  three 
groups — electrolytes,  half -electrolytes  and  non-electrolytes.  The  term 
"electrolyte"  has  been  frequently  used  in  two  different  senses — some- 
times referring  to  the  solute  and  at  other  times  used  to  designate  the  con- 
ducting solution.  It  is  more  properly  used  in  this  latter  sense.  The 
word  "  iongen  "  has  been  proposed  as  a  suitable  name  for  those  substances 
which  dissociate  when  placed  in  solution.  This  seems  to  be  more  consis- 
tent with  facts,  as  salts,  acid  and  bases,  with  the  exception  of  fused  salts, 
are  in  themselves  non-conductors  of  the  electric  current.  Solutions  of 
salts  and  strong  acids  and  bases  are  included  in  the  class  known  as  elec- 
trolytes; solutions  of  such  weak  acids  and  bases,  as  acetic  acid  and  ammo- 
nium hydroxide,  which  are  poor  conductors  of  the  electric  current,  are 
grouped  as  half-electrolytes,  while  all  other  substances,  largely  organic, 
which  do  not  conduct  are  called  non-electrolytes. 

Before  proceeding  with  special  phases  and  applications  of  the  theory, 
it  might  be  well  to  consider  a  number  of  explanations  which  have  been 
volunteered  in  an  effort  to  answer  the  question  as  to  the  source  of  the  elec- 
tric charges  possessed  by  the  ions. 

Source  of  Electric  Charges. — Le  Blanc  thinks  that  in  a  compound  like 


132  EXPERIMENTAL  CHEMISTRY. 

potassium  iodide  there  remains  some  chemical  energy  after  the  metallic 
potassium  and  iodine  have  chemically  combined,  and  that  this  chemical 
energy,  through  the  influence  of  the  solvent,  is  transformed  into  electrical 
energy  which  is  seated  in  the  charges  of  the  ions.  By  the  aid  of  an 
electric  current  during  electrolysis  it  is  possible  to  add  to  these  ions  the 
energy  which  they  originally  possessed  as  neutral  substances,  when  they 
will  separate  in  the  ordinary  molecular  forms  at  the  electrodes. 

According  to  Nernst,  the  dielectric  constant  of  a  liquid  and  its  dissociat- 
ing power  are  in  direct  proportion.  He  holds  that  the  forces  which  bind 
the  ions  together  to  form  a  molecule  are  due  to  electrical  attractions  be- 
tween the  oppositely  charged  ions.  The  specific  inductive  capacity  of 
water  at  20°  C.  is  81.1  times  that  of  air;  therefore,  since  the  ions  are  sup- 
posed to  have  a  constant  charge,  the  force  exerted  between  the  ions  in  a 
molecule,  tending  to  prevent  the  dissociation  will  be  81.1  times  less  in 
water  than  in  air.  The  dielectric  constant  of  ethyl  alcohol  is  26,  for 
ethyl  ether  4.36,  and  for  carbon  disulphide  2.6. 

The  "electron  theory,"  as  advanced  by  J.  J.  Thomson,  suggests"that 
there  is  only  one  kind  of  ultimate  particle — the  electron,  an  atom  of  elec- 
tricity which  has  a  mass  af  about  one-thousandth  that  of  a  hydrogen  atom. 
This  theory  necessitates  a  revision  of  our  former  notions  of  the  chemical 
atom.  It  assumes  "that  atoms  of  the  various  elements  are  collections 
(constellations)  of  these  electrons — positive  and  negative  charges  held 
together  mainly  by  their  electric  attractions,  associated  with  more  or  less 
ether."  These  corpuscles  are  conceived  of  as  being  in  a  state  of  rapid 
vibration.  In  some  atoms  the  velocities  of  the  electrons  (negative  charges) 
may  be  so  great  that  a  corpuscle  escapes  at  once  from  the  atom,  thus 
leaving  the  atom  positively  charged;  or  the  reverse  may  be  the  case,  leaving 
the  atom  negatively  charged.  A  negatively  charged  body  contains  an 
excess  of  electrons;  a  positively  charged  one,  a  deject.  It  is  thought,  for 
example,  that  when  sodium  and  chlorine,  each  representing  a  particular 
"constellation"  of  electrons,  combine  to  form  sodium  chloride  that  the 
two  groups  of  electrons  are  modified.  Now,  it  is  difficult  to  break  up  a 
molecule  into  atoms  out  of  which  it  was  formed,  but  comparatively  easy 
to  split  it  in  a  slightly  different  manner,  so  as  to  leave  one  electron  in 
excess  with  the  chlorine,  and  a  deject  of  one  with  the  sodium,  with  the  re- 
sult that  they  become  charged  positively  and  negatively,  "respectively. 
The  organic  chemist  is  familiar  with  the  fact  that  it  may  be  quite  difficult 
to  split  a  long  carbon  compound  at  a  particular  bond  or  link,  but  very 
easy  perhaps  to  break  it  at  some  other  link.  It  follows  from  the  above 
that  "since  electrons  possess  a  definite  though  very  small  mass,  the  asso- 
ciation with,  or  removal  of,  the  electric  charges  from  matter  causes  a  differ- 
ence in  its  mass. 

d  +  e  _  Q' 

Na  — e  —  Na 

+ 

Chlorine  has  an  atomic  mass  of  35.45;  therefore,  the  mass  of  a  chlorine 
ion  would  be  35.451,  and  the  weight  of  a  sodium  ion  would  be  .001  less 


NOTE    ON    THE    MODERN    THEORY    OF    SOLUTION.  133 

than  the  atomic  weight.  Here  as  elsewhere  the  student  should  not  con- 
fuse fact  and  theory. 

In  an  article  by  G.  T.  Beilby,  which  recently  appeared  in  The  Chemi- 
cal News,  London,  some  very  interesting  suggestions  are  made  by  the 
author  in  regard  to  the  probable  source  of  the  energy  of  solution.  The 
following  are  extracts  from  said  article:  "  The  solute  molecules  in  a  dilute 
solution  of  any  non-volatile  solid,  are  solid  molecules  sparsly  distributed 
among  a  multitude  of  intensely  active  solvent  molecules,  the  temperature 
of  the  solution  being  frequently  many  hundred  degrees  below  that  at 
which  they  could  of  themselves  assume  the  greater  freedom  of  the  liquid 
or  gaseous  state.  These  solute  molecules  have  to  a  great  extent  been  set 
free  from  the  constraining  effect  of  their  cohesive  forces,  but  it  is  important 
to  remember  that  this  freedom  has  not  been  attained  by  the  increase  of  their 
own  kinetic  energy  as  in  liquefaction  by  heat.  Their  freedom  and  the 
extra  kinetic  energy  they  have  acquired  have  in  some  way  been  imparted 
to  them  by  the  more  active  solvent  molecules;  for,  if  the  solvent  could  be 
suddenly  removed,  leaving  the  solute  molecules  still  similarly  distributed 
in  a  vacuous  space,  they  would  eventually  condense  into  a  solid  aggregate. 
This  must  be  the  case,  for  the  non-volatile  solute  has  no  measurable  vapor 
pressure  at  the  temperature  of  the  solution.  The  kinetic  energy  of  the 
solute  molecules  is  of  itself  quite  insufficient  to  endow  them  with  the  prop- 
erties of  the  gaseous  or  even  of  the  liquid  molecule,  even  when  their 
cohesive  forces  have  been  weakened  or  overcome  by  separation. 

//  the  energy  employed  in  this  separation  is  not  intrinsic  to  the  solute 
molecule  then  it  must  in  some  way  have  been  imparted  by  the  solvent  mole- 
cules. It  therefore  becomes  important  to  compare  the  energy  endowment 
of  one  set  of  molecules  with  that  of  the  other."  *  *  *  "Taking  into  con- 
sideration not  only  this  greater  store  of  energy,  but  also  the  much  smaller 
cohesive  force  of  water  as  compared  with  the  majority  of  solid  solutes, 
there  can  be  no  doubt  that  the  active  role  in  aqueous  solutions  must  be 
assigned  to  the  solvent,  not  to  the  solute  molecules. 

This  leads  to  the  important  conclusion  that  the  energy  oj  solution,  oj 
diffusion,  and  of  osmosis  is  due,  not  to  imaginary  gaseous  energy  oj  the 
solute,  but  to  the  actual  liquid  energy  oj  the  solvent"  *  *  *  "The  rude 
mechanical  jostling  to  which  the  complex  molecule  is  subjected  will  natu- 
rally tend  to  break  it  up  into  simpler  portions  which  are  mechanically 
more  stable." 

"The  view  that  the  phenomena  of  solution  depend  on  the  relative 
kinetic  energy  of  the  solvent  and  solute  molecules  appears  to  apply  with 
special  force  to  the  phenomena  of  dissociation  in  dilute  solutions.  Under 
the  gas  theory  there  does  not  appear  to  be  any  reason  why  the  solute 
molecules  should  dissociate  into  their  ions.  So  obvious  is  this  absence  of 
any  physical  motive  that  Professor  Armstrong  has  happily  referred  to 
the  dissociation  as  "  the  suicide  of  the  molecules."  "  With  the  acceptance 
of  the  view  that  phenomena  of  solution  are  largely  due  to  kinetic  energy 
of  the  solvent  molecules,  the  phenomena  of  dissociation  also  appear  to 
take  their  place  as  a  natural  result  of  this  activity,  for  consider  the  situa- 


134  EXPERIMENTAL    CHEMISTRY. 

tion  of  an  isolated  molecule  (solute),  closely  surrounded  by  and  at  the 
mercy  of  some  millions  of  water  molecules  all  in  a  state  of  intense  activity." 

"An  ideally  perfect  solution — that  is,  a  solution  of  which  the  physical 
properties  are  determined  solely  by  the  number  of  molecules  it  contains 
in  a  given  volume — must  consist  of  a  solvent  and  a  solute  which  have  no 
chemical  affinity  for  each  other,  so  that  their  molecules  will  neither  asso- 
ciate not  dissociate  in  solution.  Probably  comparatively  few  solutions 
will  be  found  which  even  approximate  to  this  ideal  perfection." 

It  appears  as  though  the  two  latter  theories  might  be  correlated  so  as 
to  become  supplementary. 

References. 

Theoretical  Chemistry,  Nernst.  Inorganic  Chemistry,  Ostwald. 

Physical  Chemistry,  Walker.  Lehrbuch    d.  Allg.  Chem.,  Ostwald. 

Physical  Chemistry,  Jones.  Scientific  Foundations,  Ostwald. 

Electro-Chemistry,  Lehfeldt.  Scientific  Memoirs,  containing  orig- 

Electro-Chemistry,  Le  Blanc.  inal  papers  of  Pfeffer,  Van't  Hoff, 

Theory  of  Electrolytic  Dissocia-  Raoult,  Arrhenius.  Edited  by  Ames. 
tion,  H.  C.  Jones. 


APPLICATION   OF   THEORY. 
Electrolysis  and  Electrical  Equivalents. 
Experiment  I.— Electrolysis. 

(a)  Refer  to  experiment  in  which  the  electrolysis  of  H2O  was  performed 
through  the  agency  of  Na2SO4.     Interpret  the  phenomena  observed  in 
the  light  of  the  "ion  theory."     Represent  the  electrolytic  reactions  by 
means  of  equations  in  which  are  shown  the  ions.     Explain  the  formation 
of  free  elements  at  the  electrodes. 

(b)  Perform  the  electrolysis  of  HC1.     Explain  mechanism. 

Careful  measurements  have  shown  that  96,580*  coulombs  (one  "Fara- 
day") of  electricity  will  deposit  one  gram-equivalent  of  any  substance. 
This  quantity  of  electricity  is  known  as  the  electro-chemical  constant.  A 
short  rule  for  determining  the  gram-equivalent  is:  divide  the  atomic 
weight  of  an  element,  or  the  sum  of  the  atomic  weights  if  it  is  a  radical, 
by  its  valency;  the  quotient  will  represent  the  number  of  grams  of  that 
substance  deposited  by  96,580  coulombs.  The  electro-chemical  equivalent 
may  be  found  by  dividing  the  gram-equivalent  by  96,580.  It  is  the  amount 
deposited  by  one  coulomb.  Example:  Atomic  weight  of  oxygen  is  16; 
its  valency  is  2;  therefore  8  grams  is  the  gram-equivalent.  For  silver  it 
is,  107.93  "*"  I  =  IO7-93  grams.  The  electro-chemical  equivalent  of  silver 
is,  107.92  -v-  96.580  =  .0011175  grams.  Recall  Faraday's  laws. 
*  Richards',  T.  W.,  value. 


NOTE    ON   THE   MODERN    THEORY    OF    SOLUTION.  135 

Names   of   Ions. — The  following  list  is  based  on  Walker's  system: 

Anions.  Cations. 

Cl'  Chloridion  Na'  Natrion  (Sodium) 

CIO/  Chloranion  K'  Kalion 

CIO/  Perchloranion  NH/  Ammonion 

NO/  Nitranion  Ca"  Calcion 

SO/'  Sulphanion  H'  Hydrion 

SO3"  Sulphosion  Fe'"  Triferrion 

S"  Sulphidion  Fe"  Diferrion 

HC2H3O/  Acetanion  Cu"  Dicuprion 

OH'  Hydroxidion  Ag'  Argention 

Some  of  the  most  interesting  things  which  are  learned  in  regard  to  ions 
by  means  of  experiments  in  electrolysis  are,  (a)  no  electrical  energy  is 
consumed  in  the  production  of  ions;  their  birth  is  antecendent  to  the  pass- 
ing of  the  current,  (b)  they  carry  electrical  charges  which  are  proportional 
to  their  relative  valencies,  (c)  their  relative  migration  velocities  differ 
greatly,  but  the  migration  velocity  of  a  particular  ion  is  independent  of  the 
nature  of  the  co-existent  ions,  and  (d)  the  absolute  velocity  of  the  most 
speedy  ion  is  very  slow.  The  following  absolute  velocities  at  18°  calcu- 
lated for  infinitely  diluted  aqueous  solutions  are  given  by  Kohlrausch. 
The  difference  of  potential  between  the  electrodes  i  cm.  apart  was  i  volt. 

K       =   .00066  cm.  H        =  .00320  cm. 

NH4  =  .00066  cm.  Cl       =   .00069  cm. 

Na     =   .00045  cm.  NO3  =   .00064  cm. 

Li       =  .00036  cm.  C1O3  =  ,00057  cm- 

Ag      =  .00057  cm-  OH     =   .00181  cm. 

It  is  seen  that  the  movement  of  the  ions  through  practically  pure  water 
is  very  slow.  Walker  calculates  the  force"  required  to  drive  i  gram  of 
hydrion  through  water  at  the  rate  of  i  cm.  per  second  to  be  equal  to  about 
320,000  tons  weight.  It  is  obvious  that  the  ions  experience  great  resist- 
ance to  their  movements.  This  is  said  to  be  due,  at  least  in  part,  to  the 
"  hydration  of  the  ion,"  i.e.,  the  comparatively  large  amount  of  water  which 
travels  with  the  ions. 

CONDUCTIVITIES. 

Experiment  II.  (L.  T.)  Electrolytes,  Half -Electrolytes  and  Non-Elec- 
trolytes. 

(Instructions.)  By  means  of  a  number  of  conductivity  cells  (Figs. 
26  and  27)  and  an  alternating  current  suitably  applied,  determine  the 
relative  electrical  conductivities  of  the  following  substances.  On  basis 
of  results  arrange  substances  under  the  three  heads  suggested  above. 
Sugar  (C12H22OU)  solution,  toluene  (C7Hg),  NaCl  solution,  C2H5OH, 
chloroform  (CHC13),  HC1  solution,  NaOH  solution,  acetic  acid  (HC2H3O2) 
solution,  NH4OH,  distilled  H2O,  dilute  solution  of  C2H5OH,  NH4C1  solu- 
tion, HNO3  solution,  KC2H3O2  solution,  C6H6. 


i36 


EXPERIMENTAL  CHEMISTRY. 


Note. — It  is  suggested  that  .iN  or  .5N  solutions  of  acids,  bases  and  salts 
be  used  in  above  experiment. 

In  view  of  your  experiments,  what  would  you  say  of  the  conductivity 
of  organic  compounds  as  compared  with  inorganic  substances  in  aqueous 
solution  ? 

There  is,  accurately  speaking,  no  sharp  line  of  demarcation  between  these 
classes.  The  distinction  is  based,  however,  on  degree  of  conductivity. 
It  should  be  recalled  that  electrolytes  produced  abnormal  osmotic  press- 


FIG.  26. 


FIG.  27. 


ures,  abnormal  depressions  of  the  freezing  point  and  abnormal  lower- 
ings  of  the  vapor  pressure;  that  is,  they  (solutes)  apparently  dissociated  in 
water  in  a  manner  similar  to  the  dissociation  of  gases  under  favorable 
conditions  of  pressure  and  temperature. 

N204  <=>  N02  +  N02 


On  the  other  hand,  the  non-electrolytes  yield  normal  results  relative  to 
osmotic  pressure,  depression  of  freezing  point,  and  lowering  of  vapor 
pressure.  It  seems,  therefore,  that  conductivity  is  due  to  these  little 
particles  into  which  the  molecules  are  dissociated,  namely,  anions  and 
cations.  And  because  these  ions  are  attracted  to  the  positive  and  nega- 
tive electrodes  respectively  of  an  electrolytic  cell,  they  must  be  oppositely 
charged.  This  kind  of  splitting  up  of  the  molecule  has  been  called 
"electrolytic  dissociation"  to  distinguish  it  from  gaseous  dissociation 
where  the  particles  are  not  apparently  electrically  charged. 

Other  conditions  being  the  same,  conductivity  depends  upon  the  degree 
of  ionization,  i.e.,  the  number  of  ions  present. 

Experiment  III.  —  Ionization  in  Solution. 

(a)  Nature  of  Solute.  Determine  relative  electrical  conductivity 
of  aqueous  solutions  of  the  following:  NaCl,  C2H5OH,  HNO3,  CHC13, 


NOTE    ON   THE    MODERN   THEORY    OF    SOLUTION.  137 

KOH,  NH4OH,  KNO3,  NaOH,  Na,,SO4,  Ca(OH)2,  HC1,  HC2H3O2  and 
HgCl2.  Tabulate  data  under  the  following  heads:  acids,  bases,  salts, 
organic  compounds,  so  that  substances  .having  greatest  conductivities 
will  stand  at  the  head  of  their  respective  columns. 

Is  it  evident  from  above  that  the  ionization  of  different  solutes  vary 
greatly  ?  Write  the  formula  for  one  substance  in  each  column  showing 
the  ions  formed  by  its  dissociation.  What  ions  are  common  to  all  acids  ? 
Bases?  If  the  "strength"  of  acids  and  bases  depends  upon  the  relative 
concentrations  of  hydrogen  and  hydroxidions  respectively,  which  of  the 
above  represent  the  "  strong  "  acids  and  bases  ?  Indicate  by  a  star  placed 
opposite  the  formula. 

As  a  general  rule,  it  may  be  stated  that  aqueous  solutions  of  salts  and 
the  so-called  "strong"  acids  and  bases  are  good  conductors  because  they 
are  largely  ionized.  Mercuric  chloride  is  an  exception,  but  not  the  only 
one.  Dissociation  of  a  substance  may  be  said  to  be  one  of  its  specific 
properties,  and,  as  such,  it  would  be  expected  that  different  substances 
possess  this  property  in  varying  degrees.  The  beginner  is  again  cautioned 
against  concluding  that  because  a  substance  is  very  soluble  that  it  is  very 
largely  dissociated.  Sugar  is  quite  soluble,  but  we  have  no  evidence  that 
it  dissociates. 

(b)  (i)  Effect  and  nature  of  solvent.  Proceed  as  in  foregoing  ex- 
periments to  ascertain  the  relative  conductivity  of:  (i)  distilled  H2O,  (2) 
dry  NaCl,  (3)  dry  C12H22On,  (4)  dry  Nal,  (5)  aqueous  solutions  of  (2), 
(3),  (4).  Using  C2H5OH,  CHC13,  and  C7H8  instead  of  water  repeat  (5). 

Add  alcohol  to  a- NaCl  solution  from  time  to  time,  and  as  frequently  note 
effect  on  the  conductivity  of  the  solution.  Results  ? 

(b)  (2)  Prepare  some  perfectly  dry  HC1  gas  by  bubbling  it  through 
concentrated  H2SO4;  run  the  gas  into  toluene  (C7H8)  or  benzene  (C6H6) 
until  latter  is  saturated,  then  run  gas  into  distilled  H2O;  test  each  of  the 
solutions  with  litmus  paper.  Results  ?  Now  place  the  solutions  of  HC1 
in  two  separate  and  thoroughly  cleaned  and  dried  conductivity  cells  and 
determine  relative  conductivity.  Results  ?  Does  the  toluene  solution  of 
HC1  show  any  acid  properties  whatever  ?  Are  there  any  ions  in  this  solu- 
tion ?  Your  reasons  ?  Set  aside  this  latter  solution  for  future  use. 

What  would  you  say  of  the  dissociating  power  of  H2O  as  compared 
with  other  solvents  tested  ? 

The  nature  of  the  solvent  plays  a  very  important  role  in  determining 
whether  a  resulting  solution  will  conduct  or  not,  i.e.,  whether  the  solute 
will  or  will  not  ionize.  Therefore  conductivity  does  not  depend  alone  on 
either  solute  or  solvent.  Water  is  the  most  efficient  dissociating  agent 
with  which  we  are  acquainted. 

At  ordinary  temperatures  pure  substances  appear  to  have  small  con- 
ductivities, providing  the  transfer  is  electrolytic.  At  higher  temperatures 
fused  salts  and  metallic  oxides  are  good  electrolytic  conductors.  It  was 
maintained  at  one  time  that  only  aqueous  solutions  were  good  electrolytes, 
but  it  has  since  been  learned  that  certain  non-aqueous  solutions  frequently 
give  higher  conductivities.  Liquefied  NH3,  SO2  and  HCN  dissolve 


138  EXPERIMENTAL  CHEMISTRY. 

salts  and  make  good  conductors.  Their  properties,  however,  are  much 
more  complex  than  those  of  water. 

A  satisfactory  and  final  explanation  as  to  why  water  or  any  other  solvent 
acts  in  such  a  manner  has  not  been  volunteered.  It  has  been  suggested 
as  a  general  rule: 

"  When  the  solvent  is  associated,  the  dissolved  molecules  are  mostly 
simple,  and  vice  versa." 

There  is  experimental  evidence  to  support  the  idea  that  in  pure  water 
there  is  found  in  addition  to  H2O  <=±  H*  +  OH',  polymerized  or  asso- 
ciated molecules  of  water  with  formulae  probably  varying  from  H4O2 — 
H8O4.  Water  itself  when  pure  conducts  only  in  a  minimum  degree. 
It  should  be  kept  in  mind  that  the  hydrions  and  hydroxidions  of  dissoci- 
ated water  may  unite  with  other  oppositely  charged  ions,  and  the  com- 
pound thus  formed  may  or  may  not  dissociate  depending  upon  its  nature 
and  the  solvent. 

It  has  been  calculated  that  ten  to  eleven  million  liters  of  pure  water 
contain  one  gram  of  hydrions  and  seventeen  grams  of  hydroxidions. 

When  NaCl  is  dissolved  in  water  a  portion  of  it  dissociates  as 
per  the  equation,  NaCl  <=±  Na'  +  Cl',  but  owing  to 

H2O  «=±  O  H'  +  H- 

JI         Jf 
NaOH  HC1, 

the  dissociation  of  water  the  above  equations  represent  the  reactions  that 
will  undoubtedly  occur,  as  well  as  the  number  of  different  particles  present. 

(c)  Effect  of  dilution.  Test  the  conductivity  of  glacial  acetic  acid  or 
concentrated  sulphuric  acid.  Add  a  few  drops  of  water;  test.  Repeat 
operations  many  times.  Record  all  results.  Is  ionization  increased  or 
diminished  by  dilution?  Write  equations  showing  the  dissociation  on 
addition  of  water  to  HC2H3O2.  How  many  particles  are  present  in  the 
solution?  Write  similar  equations  for  NaOH  and  NajSO^ 

It  is  quite  obvious  that  if  increased  conductivity  is  due  to  an  increase  in 
ionization  in  the  above  experiment  that  dilution  produces  a  greater  degree 
of  dissociation. 

Arrhenius  established  the  following  law  which  although  not  absolutely 
correct  is  sufficiently  accurate  for  practical  purposes. 

"  The  degree  of  dissociation  of  a  substance  in  a  solution  is  equal  to 
the  ratio  of  the  equivalent  conductivity  of  that  solution  to  its  equivalent 
conductivity  at  infinite  dilution." 

Av 


Another  law  of  much  importance  to  the  chemist  is  known  as  Ostwald's 
Law  of  Dilution.  It  applys  to  a  binary  electrolyte,  i.e.,  when  each  mole- 
cule forms  two  ions.  The  law  involves  the  mass  law. 


NOTE    ON   THE    MODERN    THEORY    OF    SOLUTION.  139 

"  The  product  of  the  concentration  of  the  two  ions  divided  by  the  con- 
centration of  the  undissociated  part  is  a  constant" 
This  law  may  be  represented  in  formula  as  follows: 

(1)  Na'  +  Cl'^ 

(2)  C3  .  C2  =  K 

C2.C3 

(3)  K  — 


KG, 

is  known  as  the  "dissociation  constant"  and  its  value  varies  with  each 
solute  and  solvent  and  temperature.  It  states,  however,  a  constant  re- 
lation which  exists  between  the  ions  and  the  undissociated  molecules  of 
an  electrolyte.  The  relation  of  dilution  to  degree  of  ionization  may  be 
shown  by  altering  the  form  of  the  equation.  Let  y  equal  the  fraction 
that  is  dissociated,  then  i  —  y  equals  the  part  which  is  in  the  molecular 
condition,  and  v  the  volume  to  which  the  solution  is  diluted. 

y  y  i  —  y 

(4)     Cs  =  -  ;  C2  =  -  ;  Q  =  --  . 


V  V  V 

y2 

y2 


(s)     '  -  y 


—  =  K;  (i  —  y)  v  =  K. 


It  is  evident  on  mere  inspection  that  if  we  dilute  the  electrolyte,  i.e., 
increase  its  volume  (v)  by  addition  of  the  pure  solvent,  then  we  diminish 
all  the  concentrations  of  the  substances  represented  in  the  formula — 
see  (3);  but  this  will  affect  the  numerator  more  than  the  denominator 
because  the  former  is  a  product  of  two  concentrations.  If  such  were  true 
then  K  would  not  remain  constant,  which  is  contrary  to  our  experimental 
data.  In  order  then  for  the  value  of  K  to  remain  constant  the  concentra- 
tions of  C2  and  C3  must  become  greater  through  the  dissociation  of  the 
undissociated  molecules  represented  by  Q.  The  formula  therefore  rep- 
resents that  dilution  produces  a  greater  degree  of  ionization. 

Although  the  law  of  dissociation  as  given  by  Ostwald  does  not  apply  to 
strong  electrolytes,  yet  the  formula  may  be  taken  as  practically  correct  up 
to  a  concentration  of  .oiN. 

The  dissociation  constant  K  is  a  characteristic  of  every  compound. 
If  the  degree  of  ionization  is  large  for  any  substance  then  the  value  of 
K  will  be  large  and  vice  versa.  We  shall  see  later  how  this  "constant" 
is  really  a  measure  of  the  strength  of  acids. 

The  form  of  the  law  of  dilution  suggests  a  condition  of  equilibrium  as 


140  EXPERIMENTAL  CHEMISTRY. 

existing  between  the  product  of  the  ions  on  one  side  and  the  molecules  on 
the  other.  As  a  matter  of  fact  this  is  the  real  statement  of  the  relations 
existent.  The  ions  are  in  equilibrium  with  the  undissociated  molecules. 
If  the  concentration  of  any  of  the  particles  is  altered,  say  C2  is  increased, 
then  some  of  C3  will  combine  with  C2  to  increase  the  value  of  Cx  until  the 
mathematical  form  of  the  equation  is  re-established.  Again,  if  the  value  of 
Cx  should  be  increased,  the  reaction  would  run  in  such  a  direction  as  to 
increase  the  values  of  C2  and  C3  until  original  conditions  prevail.  Such 
changes  are  called  reversible — and  the  change  itself  is  known  as  a  balanced 
action. 

The  effect  of  temperature  and  pressure  upon  ionization  is  easily  antici- 
pated by  applying  Le  Chatelier's  Theorem.  In  some  cases  the  heat  of 
dissociation  is  positive  while  in  others  it  is  negative. 

COLOR    OF    IONS. 

The  theory  of  electrolytic  dissociation  requires  that  the  color  of  an  elec- 
trolyte, i.e.,  a  dilute  salt  solution  shall  depend  upon  the  color  of  its  free 
ions.  There  is  much  experimental  evidence  to  support  this  view. 

Experiment  IV. — Persistency  of  the  Color  of  an  Ion. 

(a)  Observe  the  colors  of  solutions  of  the  following  substances  as 
you  find  them  among  the  "  shelf  reagents."     Write  the  equations  showing 
the  ions  formed  by  dissociation.     Tabulate  results,  stating  the  color  of 
the  respective  anions  and  cations: 

NaCl,  NaBr,  KNO3,  KC1,  Na2SO4,  KC2H3O2,  NH4C1,  KI,  BaCl2, 
CaCl2,  KBr,  NH4C2H3O2. 

(b)  Now  examine  solutions  of  Cu(C2H3O2)2,  CuCl2,  Cu(NO3)2,  CuSO4; 
compare  the  color  of  the  solutions  with  the  substance  in  the  dry  solid 
state;  write  equations  showing  ions  formed  by  dissociation;  by  means  of 
(a)  tabulate  colors  of  anions;  what  ion  is  the  cation  common  to  all  the 
solutions  ?     What  is  the  probable  source  of  the  color  common  to  the  solu- 
tions?    Is  the  color  likely  due  to  the  salt  in  the  molecular  condition? 
Why  ?     What  is  the  color  of  anhydrous  copper  sulphate  ?     (Recall  a  for- 
mer experiment.)     Note  colors  of  solutions  of  Co(NO3)2,  CoCl2,  K2CrO4. 
State  source  of  color  and  your  reasons  for  thinking  so. 

(c)  Place  a  small  pinch  of   CuBr2  in  a  test  tube.     What  is  the  color 
of  the  dry  salt  ?     Add  a  few  drops  of  water.     Observe  color.     Add  water 
to  test  tube  until  one-fourth  full.     Is  it  probable  that  the  dominant  color 
is   due   to   molecular    CuBr2?     Your  reasons?     Now  fill  the  test  tube 
with  water.     What  is  the  color  ?     Is  it  similar  in  color  to  solutions  of  cop- 
per salts  used  in  (b)  ?     Why  was  it  necessary  to  dilute  the  solution  to  so 
great  a  degree  to  procure  final  color?     Write  the  formula  for  Ostwald's 
Law  of  Dilution.     Interpret  above  experiment  in  terms  of  it. 

(d)  Effect  of  magnitude  of  electrical  charge  on  color  of  ion.     Prepare 
dilute  solutions  of  FeSO4,  and  FeCl3.     Show  by  means  of  formulae  the  ions 


NOTE    ON    THE    MODERN    THEORY    OF    SOLUTION.  141 

formed  in  the  respective  solutions.  What  is  the  color  of  the  anions  in 
both?  Is  the  iron  ion  common  to  both  solutions?  Is  its  electrical 
charge  (valence)  identical  in  both  cases  ?  Is  its  energy  content  identical 
in  both  solutions  ?  Would  you  infer  that  the  latter  had  something  to  do 
with  its  properties?  What  are  your  conclusions  in  regard  to  the  differ- 
ence in  color  of  the  two  solutions  ?  Are  the  two  iron  ions  identical  ?  In  view 
of  (a),  (b),  (c),  would  you  consider  it  an  extravagant  assumption  to  assume 
that  any  particular  ion  always  maintains  its  identity,  i.e.,  possesses  an  in- 
dividual set  of  chemical  and  physical  properties? 

Ostwald  in  particular  has  studied  this  question  of  the  absorption  spectra 
of  solutions  from  the  stand-point  of  the  dissociation  theory,  and  the  fol- 
lowing conclusion  was  reached: 

"That  the  spectra  of  dilute  solutions  of  salts  containing  the  same 
colored  ion  are  identical." 

CHEMICAL    CONDUCT    OF    IONS. 

With  comparatively  few  exceptions  chemical  reactions  are  the  result 
of  the  mutual  interaction  of  ions. 

Experiment  V. — Ions  Necessary  to  Chemical  Reaction — Acids. 

(a)  Place  a  portion  of  the  concentrated  H2SO4  or  HC2H3O2,  which 
would  not  conduct  the  electric  current,  upon  a  piece  of  dry  zinc.     Results  ? 
Dilute  the  acids  with  four  or  five  times  their  volume  of  distilled  water. 
Results?     Explain   and  write   equations  showing  action  of  water  upon 
acids  and  of  solutions  upon  zinc. 

(b)  Pour  some  of  the  toluene  saturated  with  dry  HC1  upon  a  dry  piece 
of  marble.     Results?     Did  the  above  solution  act  like  an  electrolyte? 
Add  a  little  water  to  a  portion  of  the  toluene  solution  of  HC1;  test  its 
conductivity;  test  its  action  upon  a  dry  piece  of  marble.     Results?     Ex- 
plain.    Equations  ? 

(c)  Recall   or   repeat   the   experiment   of    mixing   dry  NaHCO3  and 
H2C4H4O6  in  a  mortar.    Observe  whether  any  chemical  action  takes  place. 
Add  water  to  the  mixture.     Results  ?     Explain  in  terms  of  the  dissocia- 
tion theory.     Equations? 

Concentrated  acids  are  usually  very  slightly  ionized;  that  is,  they  yield 
but  few  ions,  their  conductivity  is  low  and  their  chemical  activity  as  an 
acid  is  correspondingly  weak.  Acids  are  substances  which  when  dis- 
solved in  water  or  other  dissociating  solvents,  yield  hydrions.  This  hydrion 
is  common  to  all  acids  and  is  the  source  of  the  acidic  properties  of  the 
general  class  of  substances  known  as  acids.  The  hydrion  is  a  colorless 
substance  composed  of  one  atom  of  hydrogen  bearing  one  electrical 
charge.  It  has  a  sour,  acid  taste  and  turns  blue  litmus  red.  It  may 
transfer  its  charge  to  some  metals  like  zinc  and  magnesium  and  suffer 
displacement  from  the  solution.  Zn  +  2H'  +  SO"4  —  Zn"  +-SO"4  + 


142  EXPERIMENTAL  CHEMISTRY. 

Hj.  It  combines  with  the  hydroxidion  to  form  water.  Its  migration  ve- 
locity is  much  greater  than  that  of  any  of  the  other  common  ions,  there- 
fore it  confers  large  conductivities  upon  those  solutions  in  which  it  is  con- 
tained, providing  there  is  sufficient  concentration  of  said  ions. 

It  follows  that  the  "strength"  of  an  acid  will  depend  upon  the  con- 
centration of  hydrions  which  it  can  supply.  In  other  words,  those  acids 
whose  "dissociation  constants"  (K)  are  large,  are  known  as  "strong" 
acids  and  good  conductors,  and  vice  versa.  Conductivity  measurements 
are  approximate  measurements  of  the  relative  "strength"  of  acids. 

Experiment  VI.— "Strength"  of  Acids. 

Determine  the  conductivities  of  the  following  acids  in  the  concentrated 
and  in  the  diluted  conditions:  H3PO4,  H2SO4,  HC1,  HC2H3O2,  HNO3. 
Tabulate  results,  arranging  acids  in  order  of  increasing  conductivity. 

Any  process  which  removes  hydrions  from  a  solution  or  prevents  their 
appearance  in  same  will  eliminate  all  acid  properties.  Toluene  prevented 
the  dissociation  of  HC1,  hence  neither  conductivity  nor  acid  reaction  of 
the  toluene  solution  of  hydrogen  chloride. 

Experiment  VII.— "  Strength  "  Of  Bases. 

Proceed  as  in  former  experiments  to  ascertain  the  relative  conductivities 
of  iN.  solutions  of  NH4OH,  NaOH,  Ca(OH)2,  and  KOH.  Tabulate  re- 
sults. Write  an  equation  for  each  substance  showing  nature  of  dissociation. 
Which  ion  is  common  to  all  the  solutions  ? 

As  in  the  case  of  acids,  we  find  solutions  of  bases  giving  varying  con- 
ductivities, depending  largely  upon  their  relative  degree  of  ionization. 
Bases  are  substances,  the  aqueous  solutions  of  which  yield  hydroxidions.  It 
is  this  ion  which  confers  the  basic  properties  upon  the  class  of  substances 
known  as  bases.  The  hydroxidion  is  a  colorless  substance  composed  of 
one  atom  each  of  hydrogen  and  oxygen  (hydroxyl),  bearing  one  electrical 
charge.  It  possesses  a  caustic,  lye-like  taste  and  turns  red  litmus  blue. 
It  combines  with  the  hydrion  to  form  water.  Its  migration  velocity, 
though  much  greater  than  other  ions,  is  but  little  more  than  one-half  the 
velocity  of  the  hydrogen  ion.  The  "strength"  of  a  base  will  depend 
upon  its  "dissociation  constant."  "Strong"  bases  will  have  a  constant 
whose  value  is  comparatively  large;  "weak"  bases,  the  reverse. 

HC1       -»  H-  +  Cl' 
HN03   —  IT  +  N0'3 
HBr       —  H-  +  Br' 
NaOH  —  OH'  +  Na! 
KOH    —  OH'  +  K- 
LiOH    —  OH'  +  Li- 


NOTE    ON    THE    MODERN   THEORY    OF    SOLUTION.  143 

Neutralization.— Salts. 

H-  +  Cl'  +  Na-  +  OH'  —  Na-  +  Cl'  +  H2O  +  13700  cals. 
H-  +  Cl'  +  K-  +  OH'    —  K-  +  Cl'  +  H2O  +  13700  cals. 
H-  +  Cl'  +  Li'  +  OH'   —  Li'  +  Cl'  +  H2O  +  13700  cals. 

H-  +  OH'    —  H2O  +  13700  cals. 

H-  +  Cl'  +  NH-4  +  OH'        —  NH-4  +  Cl'  +  H2O  +  12200  cals. 
H-  +  C2H30'2  +  Na-  +  OH'  —  Na-  +  C2H3O2  +  H2O  +  13300  cals. 
H-  +  F'  +  Na-  +  OH'  ->  Na'  +  F'  +  H2O  +  16200  cals. 

Neutralization  is  the  process  whereby  the  hydrions  of  an  acid  and 
hydroxidions  of  a  base  combine  to  form  water. 

THERMO-CHEMICAL    SUPPORT    OF    THE    DISSOCIATION    THEORY. 

The  "heat  of  neutralization"  of  strong  acids  and  bases  furnishes  one 
of  the  strongest  supports  for  the  "ion  theory."  Jones  says: 

Since  all  processes  of  neutralization  of  completely  dissociated  acids  and 
bases  are  the  same,  the  heat  of  neutralization  of  all  such  acids  and  bases 
must  be  a  constant,  and  must  be  the  heat  of  combination  of  a  gram  equiv- 
alent of  hydroxyl  and  hydrogen  ions. 

It  may  be  added,  if  either  the  acid  or  base  is  "weak,"  the  heat  of 
neutralization  will  not  be  13700  cals,  but  will  vary  from  that  value;  it  may 
be  more  or  less. 

Change  in  Volume  Support. — Another  bit  of  evidence  said  to  support  the 
theory  of  solution  is  the  uniform  contraction  in  volume  observed  during  the 
neutralization  of  strong  acids  and  bases.  When  1000  cm.3  of  a  normal  solu- 
tion of  a  strong  base  are  used  to  neutralize  1000  cm.3  of  a  strong  acid,  the 
resulting  mixture  has  a  volume  which  is  always  20  cm.3  less  than  the  sum 
of  the  two  original  volumes.  The  inference  usually  made  is — the  i  gram 
of  hydrions  and  1 7  grams  of  hydroxidions  occupy  a  larger  volume  by  20  cm.3 
than  1 8  grams  of  water  which  they  form  as  the  result  of  the  neutralization. 

A  salt  is  formed  during  the  process  of  neutralization,  and  if  the  latter 
is  complete  it  may  be  said  to  be  the  result  of  the  union  of  the  cation  of 
the  base  and  the  anion  of  the  acid. 

Salts  are  named,  as  a  rule,  according  to  composition  and  without  re- 
gard to  their  conduct.  Recall  former  classification. 

Experiment  VIII.— Salts. 

(a)  Formation  of  an  insoluble  salt.     To  a  few  cm.3  of  AgNO3  solution 
procured  from  the  "shelf  reagents,"  add  5  cm.3  of  distilled  water.     Shake 
vigorously.     Divide  into  four  equal  portions,  using  test  tubes  as  recep- 
tacles.    To  one  portion  add  a  few  drops  of  HC1,  to  another  a  little  CHC13, 
to  another  a  few  drops  of  a  NaCl  solution,  to  the  remaining  tube  add  a 
small  quantity  of  a  KC1O3  solution.     Record  all  data.     Give  equations 
showing  the  interaction  of  ions. 

(b)  Formation  of  a  complex  salt.     To  a  few  cm.3  of  AgNO3  solution 
add  two  or  three  drops  of  KCN.     Caution. — KCN  is  a  poison.     Add  a 


144  EXPERIMENTAL  CHEMISTRY. 

little  more  KCN;  continue  to  add  until  a  clear  solution  is  obtained.     Add 
a  few  drops  of  HC1.    Result  ?    Explain  various  reactions.    Give  equations. 
(c)  Repeat  above  using  a  solution  of  CuSO4'     Continue  to  add  KCN 
until  solution  is  colorless.     Explain.     Equations? 

A  "complex  salt"  is  one  which  gives  a  complex  ion,  for  example, 
KAg(CN)3  —  K-  +  Ag(CN)'2.  A  "compound  salt"  like  K2Mg(SO4)2 
manifests  less  tendency  to  form  complex  ions.  K2Mg(SO4)2  — •*  2K* 
+  Mg"  +  2SO'4.  It  would  be  in  accordance  with  facts  to  state  that 
both  "complex"  and  "compound"  salts  go  through  successive  steps  of 
ionization. 

HYDROLYSIS. 

Heretofore  we  have  spoken  of  water  as  a  neutral  substance,  but  this  is 
far  from  being  true.  It  will  be  recalled  that  wrater  dissociates  into  hy- 
drions  and  hydroxidions.  If  a  salt  when  dissolved  in  water  simply  under- 
goes dissociation  the  solution  is  neutral,  since  there  are  not  present  any 
appreciable  amounts  of  free  hydrions  and  hydroxidions.  If,  however,  the 
nature  of  the  cation  of  the  salt  is  such  that  it  forms  a  weak  base  when 
combined  with  an  hydroxyl  group,  it  will  combine  with  the  OH  ions  of  the 
water,  and  being  a  weak  base,  it  remains  to  a  great  extent  in  the  associated 
condition.  If  it  should  be  only  slightly  soluble,  it  will  be  precipitated. 
This  removal  of  cation  of  salt  and  OH  ions  of  water,  leaves  the  anion 
of  the  salt  and  the  H  ions  of  the  water  in  the  solution,  which  now  possesses 
an  acid  reaction  due  to  the  free  hydrions.  If  the  anion  of  the  salt  forms 
a  weak  acid  when  combined  with  a  hydrion,  then  by  a  similar  line  of 
reasoning  this  would  leave  free  hydroxidions  in  solution  which  would  then 
possess  an  alkaline  reaction. 

Experiment  IX. — Hydrolysis. 

(a)  Place  5  cm.3  of  distilled  water  in  a  test  tube;  test  it  with  litmus 
paper  to  assure  yourself  that  it  is  neutral.     Dissolve  a  small  crystal  of 
A1C13   or  A12(SO4)3.     Are   these   "neutral  salts"?     Test  the  solution. 
Results  ?     Explain.     Equations  ? 

(b)  Using  distilled  water  and  N^COg  repeat  above  experiment.     Re- 
sults ?     Explain.     Equations  ? 

(c)  Place  a  small  crystal  of  SbCl3  in  the  bottom  of  a  test  tube;  add 
5  cm.3  of  distilled  water;  shake  vigorously,  then  add  concentrated  HC1 
drop  by  drop  until  solution  becomes  clear  after  warming.     Add  from 
25  to  50  cm.3  of  water.     Results?     Explain.     Give  equations. 

(d)  Turn  to  the  experiment  in  your  note  book  under  "Salts,"  in  which 
you  tested  with  litmus  paper  solutions  of  various  salts.     Some  of  the 
results  were  rather  perplexing  at  the  time;  can  you  explain  away  your 
previous  difficulties? 

Hydrolysis  may  be  denned  as  a  case  of  double  decomposition  in  which 
water  is  one  of  the  chemical  reagents. 


NOTE    ON    THE    MODERN   THEORY    OF    SOLUTION. 

Per  Cent  Ionized. 
(Calculated  at  18°  C.) 


Acid. 

Per  Cent. 

Acid. 

Per  Cent. 

HC1   cone 

1  3  60 

HI    5N  25°  C. 

oo  oo 

HC1    iN 

.  .  78  .  40 

H2F2,  iN  

.    7  .00 

HNO3,  cone.  .  .  . 
HNO3,  iN  
H2SO4  cone. 

9  .  60 
82.10 
o  70 

H  HS,  .iN  
H.H2P04,  .5N  25 
HC2H3O2,  iN.  .  . 

06 
0  C....I7.00 
.40 

H2SO4    iN 

' 

51   20 

HC2H3O2  .iN 

i  40 

H.HC2O4,  .iN  25 
H  HCO3    iN   • 

°c  50.30 

O    12 

HC4H406,  iN  25 
HCN,  .iN     .      . 

0  C....   8.20 

O.OI 

Base. 

Per  cent. 

Base. 

Per  Cent. 

Li  OH,  iN  
NaOH,  iN  

63  .  oo 
.  .73.80 

Ca(OH)2,  N/64, 
Sr(OH)2  N/64,  2 

25°  C.  90.00 

<°  C.       Q3  .OO 

KOH,  iN. 

.  .  78  .00 

Ba(OH)7,  .iN,. 

.  .  7  <;  .  oo 

K  OH    .iN 

86.00 

Ba(OH)9,  N/64, 

2Z°  C.    02  .OO 

NH  OH  iN 

o  40 

NH4OH    .iN 

I  .  4O 

Salts. 

Per  Cent. 

Salts. 

Per  Cent. 

NaCl    iN 

67  .  60 

CuSO4  iN    

22  .OO 

NaCl,  ,5N  
NaCl     iN 

•  -73-40 
83  oo 

NH4C1,  iN  

Zn  SO4    iN 

74.00 
.24.00 

NaHCO3,  iN... 
NaC2H3O2,  iN.. 

52-00 
53  .00 

Zn  C12,  iN  
Hg  C12,  iN  

48  .  oo 
0.90 

Na2S04,  iN  
KC1   iN 

44-50 
75  oo 

Hg(CN)2,  iN... 
Ag  NO,,  iN.. 

Small 
.  .58.00 

KNO3,  iN  

64  .  oo 

Ca  SO4,  .oiN  .  . 

63.00 

KC2H3O2    iN 

64  oo 

BaCl2     iN 

75  .00 

K2SO4,  iN 

£2    OO 

MgSO.,  .iN.. 

45  .00 

K2CO3   iN 

40    OO 

KC1O3  -5N 

.  .  70  .00 

*  Water,  o.ooooi  per  cent,  ionized. 


IO 


146  EXPERIMENTAL  CHEMISTRY. 

Experiment  X. — Ionic  Equilibrium. 

Prepare  a  saturated  solution  of  NaCL  Add  concentrated  HC1  slowly 
drop  by  drop,  until  a  white  granular  precipitate  of  NaCl  appears.  Write 
formulae  showing  nature  of  the  dissociation  of  the  NaCl  and  the  HC1. 
Is  there  an  equilibrium  between  the  ions  and  the  undissociated  molecules 
of  a  solute?  Allow  the  salt  to  settle;  decant  the  clear  supernatant  liquid 
into  another  test  tube;  add  C2H5OH  until  a  fine  granular  precipitate  of 
NaCl  appears.  Explain. 

To  interpret  correctly  the  above  phenomena  it  is  necessary  to  recall 
certain  principles  presented  previously.  An  example  may  facilitate  the  in- 
terpretation. When  NaCl  is  dissolved  in  water  the  ions  are  in  equilibrium 
with  the  molecular  salt,  and  further,  at  a  definite  temperature  there  is  a 
constant  relation  between  the  two  as  indicated  by  the  equations  which 
follow: 

NaCl  <±  Na-  +  Cl'  HC1  <=>  H"  +  Cl' 

L-j  L,2  C3  v_x4  C5         C6. 

KQ  =  C2.C3  KC4  =  C5.C6. 

This  is  a  mathematical  expression  of  the  equilibrium.  If  the  solution 
is  saturated  then  the  right  and  left-hand  members  of  the  equation  pos- 
sess their  maximum  values,  for  K  is  constant,  i.e.,  is  independent  of  the 
concentration  of  the  solution.  If  the  solution  is  diluted  then  the  value  of 
Q  becomes  smaller.  Hence  it  follows,  that  in  a  saturated  solution  at  a 
given  temperature  for  a  given  solute  the  concentration  of  the  undisso- 
ciated molecules,  and  the  product  of  the  concentrations  of  the  ions  are 
constant.  KQ  is  called  the  "solubility  product"  and  is  a  characteristic 
of  each  substance.  C2  x  C3  is  sometimes  spoken  of  as  the  "concentration 
product"  or  "ion  product."  If  the  " solubility  product"  is  exceeded, 
supersaturation  or  precipitation  will  result. 

It  might  be  well  to  mention  in  this  connection  that  a  precipitate  is  in 
equilibrium  with  the  undissociated  molecules  remaining  in  solution. 

Na'  +  Cl'+±   NaCl  (Dslvd.) 

u 

NaCl.   (Ppt) 

Referring  again  to  the  above  experiment,  the  value  of  C3  was  greatly 
increased  by  C6,  with  the  result  that  C2,  in  order  to  preserve  the  mathe- 
matical form  of  the  equation,  became  correspondingly  small  by  associ- 
ating with  some  of  C3  and  C6  to  form  NaCl.  This  in  turn  increased  the 
value  of  Q.  When  this  latter  exceeded  the  "  solubility  product "  of  NaCl, 
salt  was  precipitated. 

SOLUTION    TENSION. 

Experiment  XI. — Displacement  of  Ions  by  a  "Free  Metal." 

(a)  Dissolve  5  grams  of  Pb(C2H3O2)2  in  100  cm.3  of  distilled  H2O; 
place  solution  in  a  small  Erlenmeyer  flask;  suspend  a  piece  of  zinc  in 


NOTE    ON    THE    MODERN    THEORY    OF    SOLUTION.  147 

the  clear  solution;  set  flask  aside  to  stand  for  several  hours.     Note  the 
displacement  of  lead  by  the  zinc. 

Pb"  +  2C2H3O'2  +  Zn->Pb  +  Zn"  +  2C2H3O/2. 

(b)  Repeat  (a)  using  CuSO4  solution  instead  of  Pb  (C2H3O2)2.  Results? 
Equation  ? 

(c)  Repeat  (b)  using  a  piece  of  clean  iron  wire  bent  into  the  form  of  a 
coil.     Results  ?     Equation  ? 

(d)  Dissolve  a  little  HgNO3  in  hot  water,  then  dip  a  coiled  piece  of 
copper  wire  into  the  solution.     Set  aside  for  an  hour.     Results?     Rub 
wire  with  piece  of  filter  paper.     Results  ?     Equation  ? 

(e)  Pour  10  cm.3  of  AgNO3  solution  into  a  crystallizing  dish;  add  a  few 
drops  of  mercury;  set  aside  for  an  hour  or  more.     Results?     Equation? 

In  all  of  the  above  experiments  were  the  displaced  ions,  anions'  or 
cations  ? 

Experiment  XII. — Displacement  of  Ions  by  a  "Free  Non-Metal." 

(a)  To  separate  solutions  of  KI  and  KBr,  or  Nal  and  NaBr,  add  a 
few  drops  of  "chlorine  water."     Results?     Add  a  little  CHC13  to  each 
test  tube  and  shake.     Results?     Explain.     Equations? 

(b)  To  a  solution  of   KI  add  a  little  "bromine  water."     Results? 
Add  a  little  CHC13  and  shake.     Results  ?     Explain.     Equations  ?     In 
the  above  experiments  were  the  displaced  ions,  cations  or  anions? 

The  inferences  which  are  drawn  as  the  result  of  the  foregoing  experi- 
ments are,  (a)  the  metals  manifest  varying  tendencies  to  press  into  solution, 
i.e.,  the  ionic  condition,  since  they  are  not  otherwise  soluble;  (b)  the  non- 
metals  relative  to  one  another,  behave  in  a  manner  very  similar  to  the 
metals.  It  is  possible  to  arrange  the  metals  in  a  series  in  the  order  of 
their  "decreasing  solution  tension." 

Alkali  metals  Lead 

Alkaline  earth  metals  Hydrogen 

Magnesium  (Arsenic) 

Aluminum  Copper 

Manganese  ,  Antimony 

Zinc  Bismuth 

Chromium  Mercury 

Cadmium  Silver 

Iron  Palladium 

Cobalt  Platinum 

Nickel  Gold 
Tin 

As  a  general  rule,  the  ions  of  any  metal  are  displaced  from  a  normal 
solution  by  any  of  the  free  metals  which  precede  it  in  the  above  series. 


148  EXPERIMENTAL  CHEMISTRY. 

The  question  undoubtedly  arises  as  to  why  metals  have  any  such 
property  as  "solution  tension."  The  following  experiments  and  discus- 
sion may  enable  the  student  to  understand  a  little  more  clearly  the  nature 
of  this  phenomena. 

Experiment  XIII. — Chemical  Energy  and  Electrical  Energy. 

(a)  Prepare  a  cold  dilute  aqueous  solution  of  H2SO4  (4  of  H2O  to  i  of 
acid).     Pour  the  solution  into  a  small  beaker.     Suspend  a  strip  of  clean 
zinc  in  the  solution;  by  means  of  a  copper  wire  connect  the  zinc  with  one 
of  the  binding  posts  of  a  galvanometer  or  a  volt-meter.     Read  volt-meter. 
Is  there  evidence  of  much  chemical  action  between  zinc  and  H2SO4? 
Bind  another  copper  wire  to  which  is  attached  a  copper  strip  to  the  other 
binding  post;  suspend  copper  strip  in  the  solution  in  such  a  manner  as  not 
to   touch   zinc   strip.     Read   volt-meter.     Results?     Is   there   more   or 
less  evidence  of  chemical  action  in  the  beaker  than  when  previously 
examined?     Observe  and  record  phenomena  manifested  by  this  typical 
cell.     Did  either  the  solution  or  the  metals  contain  electricity  originally  ? 
What  is  the  probable  source  of  the  electrical  energy?     Place  a  small 
quantity  of  "granulated"  zinc  in  the  beaker  containing  the  acid  solution 
after  removing  the  strips  of  Cu.  and  Zn. ;  test  the  escaping  gas  by  means  of 
a  lighted  match.     Results?     When  chemical  action  has  ceased,  remove 
excess  of  zinc  and  once  more  insert  the  strips  of  Zn.  and  Cu. ;  take  the 
reading  of  the  volt-meter.     Explain.     The  beaker  now  contains  a  solu- 
tion  of  what?     Evaporate  a  portion  of  the  solution  to  dryness.     Is  it 
CuSO4?     Your  reasons? 

(b)  The  above  experiment  may  be  repeated  by  substituting  for  zinc 
the   following    metals,    iron,    copper,    lead,    aluminum  and  others   as 
desired.     Arrange  the  metals  in  order  of  their  increasing  electro-motive 
force.     (E.M.F.). 

The  student  is  familiar  with  the  fact  that  when  heat  energy  is  contributed 
to  solids  or  liquids  in  sufficient  quantities,  the  molecules  pass  into  the 
space  above  the  liquid,  and  if  it  be  a  closed  system,  equilibrium  is  estab- 
lished for  a  given  temperature,  when  the  pressure  of  the  vapor  is  equal 
to  the  vapor-tension  of  the  substance.  Increase  the  amount  of  internal 
energy  of  the  substance,  i.e.,  raise  its  temperature,  and  it  is  obvious  that  its 
vapor-tension  is  also  increased. 

Every  metal  has  a  certain  solution-tension  which  tends  to  push  its 
particles  into  solution  in  the  form  of  ions.  The  metal  will  continue  to 
dissolve  until  its  solution-tension  is  in  equilibrium  with  the  solution-press- 
ure of  the  ions.  Since  energy  can  not  come  from  nothing,  it  is  said  that 
the  chemical  energy  or  the  free  energy  of  the  system  is  transformed  into 
the  electrical  charges  of  the  ions.  Metals  yield  cations,  and  non-metals — 
anions. 

Now  it  is  maintained  by  Nernst  and  others  that  the  solution-tension  of 
substances  is  analogous  to  vapor-tension.  An  application  of  the  theory 


NOTE    ON    THE    MODERN    THEORY    OF    SOLUTION.  149 

of  "electrolytic  solution-tension"  will  reveal  the  points  of  analogy  as  well 
as  the  differences. 

If  we  dip  a  piece  of  metal,  say  zinc,  into  pure  water,  then  owing  to  the 
solution-tension  of  the  metal  some  ions  will  pass  into  solution.  These 
ions  are  charged  with  positive  electricity.  Since  both  kinds  of  electricity 
must  be  simultaneously  produced  whenever  electrical  energy  comes  into 
existence,  the  metallic  zinc  becomes  negatively  electrified.  The  solution 
is  now  positively  electrified  and  the  metal  negatively  electrified.  The 
positively-charged  ions  are  attracted  by  the  negatively-charged  metal,  so 
that  a  difference  of  potential  is  established.  The  metal  will,  however, 
continue  to  press  into  solution,  thereby  acquiring  a  lower  potential,  until, 
as  mentioned  above,  the  solution-tension  is  in  equilibrium  with  the  solu- 
tion pressure  of  the  ions.  The  solution  and  the  metal  now  have  their 
greatest  difference  of  potential.  The  "single  potential"  of  metals  im- 
mersed in  pure  water  obviously  depends  upon  their  respective  solution- 
tensions.  Suppose  we  now  dip  a  metal  like  silver  into  a  solution  of  its 
ions  which  have  a  greater  ionic  solution-pressure  than  the  metal's  solution- 
tension,  this  will  result  in  the  deposition  upon  the  metal  of  some  of  its  ions. 
The  metal  will  become  positively  electrified  and  the  solution  negatively 
electrified. 

Again,  if  we  place  a  metal  in  a  solution  whose  ionic  solution-pressure 
is  just  equal  to  its  solution-tension,  equilibrium  is  established  at  once. 
Ions  will  neither  be  formed  nor  deposited,  hence  there  will  be  no  differ- 
ence of  potential  between  metal  and  solution. 

In  the  above  discussion  we  have  referred  alone  to  the  solution-tension 
of  metals  which  form  positive  ions.  Le  Blanc  says,  that  as  far  as  we  know, 
all  substances  capable  of  yielding  negative  ions  have  a  high  solution- 
tension.  This  would  leave  a  free  non-metal  like  chlorine  positively 
charged. 

We  immediately  infer  that  as  the  result  of  solution-tension,  metals  and 
non-metals  acquire  a  definite  electro-motive  force,  when  placed  in  pure 
water  or  a  solution  (usually  a  normal  solution)  containing  their  respective 
ions.  This  E.M.F.  is  frequently  called  the  " single  potential"  of  metals 
or  non-metals. 

"  The  tendency  toward  chemical  reaction  and  its  accompanying  trans- 
formation of  energy  in  a  cell  is  measured  by  the  E.  M.  F.  or  voltage.1' 
— Lehfeldt's  Electro- Chemistry. 

We  have  considered  the  questions  of  solution-tension,  the  production  of 
"single  potentials"  and  the  E.M.F.  of  a  cell,  but  we  have  not  considered 
in  detail  the  probable  source  of  energy  of  which  the  foregoing  are  mani- 
fested forms.  This  energy  is  evidently,  at  least  in  part,  derived  from  the 
chemical  energy  of  the  system.  It  was  thought  for  some  time  that  the 
whole  of  the  chemical  energy  was  converted  into  electrical  energy,  but 
this  assumption  is  by  no  means  justified  by  the  facts.  This  is  only  true 
in  the  case  of  cells  whose  E.M.F.  do  not  vary  with  temperature.  In  cer- 
tain types  of  cells  a  portion  of  the  chemical  energy  is  evolved  as  heat,  the 
cell  becoming  warmer  as  the  cell  continues  in  action;  in  other  types  the 


150  EXPERIMENTAL  CHEMISTRY. 

electrical  work  done  exceeds  the  chemical  energy  spent,  i.e.,  heat  is  ab- 
sorbed and  transformed  into  electrical  energy. 

"  Measurement  of  electro-motive  force  may  then  be  looked  upon  as  a 
means  of  determining  the  change  of  the  free  energy  in  a  chemical  reac- 
tion."— Lehfeldt's  Electro-Chemistry. 

The  E.M.F.  is  a  measure  of  the  chemical  energy  of  a  system  when  the 
process  is  reversible,  i.e.,  when  the  system  is  neither  warmed  nor  cooled 
by  the  reaction  which  produced  the  E.M.F. 

TABLE    OF    SINGLE    POTENTIALS. 

(The  free  metal  is  placed  in  a  normal  solution  of  its  ions.) 

K  =  (+  2.9)  H=-       .277 

Na  =  (+  2.54)  Cu  =  -       .606 

Ba  =  (+  2.4)  As  =  -        .62    ? 

Sr  -   (+  2.3)  Bi  =-         .67    ? 

Ca  =   +    2.28  Sb  =  -        .74   ? 

Mg  =  +   i  .21  Hg  =  —  -  i  .027 

Al  =  +     i .  oo  Ag  =  -  -  i .  048 

Mn  =  +     .80  Pd  =  -  1.07    ? 

Zn  =  +       .493  Pt  =  -  1.14   ? 

Cd  =  +      .143  Au  =  -  -  1.35   ? 

Fe  =  +       .063  I  =  +         .797 

Tl  =  +       .045  Br  =  +  1.270 

Co  =  —        .  045  O  =  +  i .  396  ? 

Ni  =  -       .049  OH  =  +  1.396? 

Sn  =  -        .07   ?  Cl  =  +  1.694 

Pb  =  -        .129  NO3  ==  +1.75 
SO4  .=  +1.90 

It  will  be  observed  that  the  relative  positions  of  the  jree  metals  in  the 
" solution- tension"  series  and  the  " E.M.F."  series  are  identical.  The 
order  in  which  the  metals  occur  in  this  series  is  especially  significant, 
inasmuch  as  it  represents  the  relative  chemical  activities  of  the  jree  metals. 
All  of  the  metals  in  the  fore  part  of  the  list  readily  oxidize  when  exposed  to 
the  air;  while  those  in  the  latter  part  of  the  list  do  not.  It  follows  from 
this  that  the  former  metals  will  be  reduced  with  difficulty  from  their  ox- 
ides, while  the  reverse  will  be  true  of  the  latter.  The  arrangement  ex- 
presses the  combining  relations  of  the  metals  with  regard  to  other  elements 
as  well  as  oxygen.  None  of  the  metals  in  the  fore  part  of  the  list  are 
found  free  in  nature.  It  should  be  remembered  in  this  connection  that  all 
of  these  same  metals  can  displace  hydrogen  from  an  acid,  while  those 
which  succeed  hydrogen  in  the  list  and  are  unable  to  displace  hydrogen 
are  found  free  in  nature.  Other  relationships  will  be  found  represented 
by  this  series. 


NOTE    ON    THE    MODERN    THEORY    OF    SOLUTION. 


HEAT    OF    IONIZATION. 


Lithium  =  +  261000  joules. 
Potassium  =  +  257000  joules. 
Sodium  =  +  237000  joules. 
Strontium  =  +  244000  joules. 
Calcium  =  +  226000  joules. 
Magnesium  =  +  225000  joules 
Aluminum  =  +  165000  joules. 
Manganese  =  +  102000  joules. 
Zinc  =  +  69500  joules. 
Cadium  =  +  34600  joules. 
Iron  (ferrous)  ==  +  42400  joules. 
Thallium  =  =  -f  3400  joules. 
Cobalt  =  +  31700  joules. 


Nickel  =  +  29400  joules. 
Tin  ==  +  2900  joules. 
Lead  =  +  2900  joules. 
Hydrogen  =  —  2300  joules. 
Copper  =    —  37200  joules. 
Mercury  (Hg')  =  —  86900  joules. 
Silver  =  —  110900  joules. 


Chlorine  =  +  167400  joules. 
Bromine  =  +  121000  joules. 
Iodine  =   +  57500  joules. 
Oxygen  (1/4  O2)  =  +  88600  joules. 


Ostwald  (Lehrbuch.  d.  allg.  Chem.)  has  calculated  the  "heat  of  ioni- 
zation"  of  one  equivalent  for  a  number  of  metals.  The  above  table  is 
quoted  from  his  work. 

"  The  heat  of  ionization  of  hydrogen  is  so  small  as  almost  to  lie  within 
the  margin  of  errors,  and  may  be  ignored.  Hence,  the  approximate  rule, 
that  the  heat  of  ionization  of  a  metal  is  practically  equal  to  its  heat  of 
solution  in  dilute  (i.e.,  completely  dissociated)  acid."— Lehfeldt's  Electro- 
chemistry. 

After  inspecting  above  table  it  is  evident  that  some  of  the  ions  possess 
more  available  energy  than  when  in  the  condition  of  the  free  metal;  other 
ions,  the  reverse. 


CHAPTER  XV. 
THE  NON-METALS,  OR  ACID-FORMING  ELEMENTS. 

THE    HALOGEN    GROUP. 

(THE  CHLORINE  FAMILY.) 
Fluorine,     F,          19 
Chlorine,    Cl,     35.45 
Bromine,  Br,     79.96 
Iodine,         I,  126.97 

These  four  elements  compose  what  is  known  as  a  "natural  group"; 
the  members  are  closely  connected  by  a  similarity  of  chemical  proper- 
ties; i.e.,  they  resemble  one  another  in  their  chemical  relations,  and 
by  combination  with  identical  substances,  produce  a  series  of  corre- 
sponding compounds  which  resemble  each  other  in  their  respective 
chemical  properties. 

The  members  of  this  group  are  usually  termed  the  halogens,  because 
of  their  tendency  to  produce  salts  resembling  sea-salt  in  their  composition. 
These  salts  are  called  haloid  salts  or  simply,  halides. 

FLUORINE,    F. 
At.  Wt.  19     Mol.  Wt.  38. 

Preparation  and  Properties. — This  element  is  not  found  free  in  nature. 
The  chief  sources  are  calcium  fluoride,  CaF2,  commonly  known  as  fluor- 
spar and  cryolite,  A1F3  3NaF.  The  latter  mineral  is  found  in  relative 
abundance  in  Greenland. 

Fluorine  is  a  pale  yellowish-green  gas  with  an  unpleasant  odor  like 
chlorine.  Although  fluorine  does  not  combine  with  oxygen  it  shows  a 
great  affinity  for  hydrogen,  with  which  it  unites  explosively,  even  in  the 
dark.  In  fact,  so  great  is  the  affinity  of  fluorine  for  hydrogen  that  many 
compounds  which  contain  the  latter  are  decomposed  when  brought  into 
contact  with  it.  It  is  recorded  that  Moissan,  who  prepared  it  in  1886, 
dropped  some  of  the  liquid  fluorine  on  the  wooden  floor  of  the  laboratory 
when  the  wood  immediately  burst  into  flame.  Because  of  its  remarkable 
chemical  activity,  it  is  exceedingly  difficult  to  prepare.  Moissan  pre- 
pared it  by  electrolysis  of  pure  liquid  H2F2  ( — 25°  C.  to—  50°  C.),  mixed 
with  a  little  KF. 

H  +F  — HF  +  37,600  cal. 


THE    NON-METALS,    OR    ACID-FORMING    ELEMENTS.  153 

Hydrogen  Derivatives. 

Experiment  I.  —  Preparation  and  Properties  of  Hydrogen  Fluorine. 

To  a  gram  of  powdered  CaF2  in  a  test  tube  add  3  cm.  3  of  concentrated 
H2SO4;  heat  the  tube  gently;  test  the  effect  of  the  evolved  gas  upon  a 
piece  of  moistened  blue  litmus  paper.  Dip  a  glass  rod  into  a  few  cm.3 
of  NH4OH  in  a  test  tube;  hold  rod  in  evolved  gas.  Results?  After 
the  test  tube  has  been  thoroughly  cleaned,  observe  the  effect  of  the  gas 
upon  the  inside  surface  of  the  glass.  Write  equations  indicating  nature 
of  above  reactions. 

Caution.  —  The  above  experiment  should  be  performed  in  the  hood. 
The  gas  must  not  come  in  contact  with  the  skin  nor  be  breathed. 

Experiment  II.  —  Etching  of  Glass  with  Hydrofluoric  Acid. 

Spread  a  smooth  thin  layer  of  paraffin  wax  upon  one  side  of  a  piece  of 
glass,  about  10  cm.  square;  allow  the  wax  to  cool  and  harden;  by  means 
of  the  sharp  end  of  a  file  or  any  pointed  instrument,  cut  through  the  wax 
to  the  glass,  making  any  chosen  design.  Cover  the  bottom  of  a  shallow 
lead  dish  with  CaF2;  add  sufficient  warm  H2SO4  to  moisten  powder  in 
dish;  place  glass  plate,  waxed  side  downward,  upon  the  lead  dish;  warm 
dish  gently  but  do  not  melt  paraffin.  After  10  min.  remove  heat,  and 
cool  dish;  remove  glass  plate  and  wash  off  the  wax.  Has  the  gas,  H2F2, 
evolved  affected  the  glass  plate  ?  Equation  ? 

Caution.  —  Use  the  greatest  care  to  avoid  breathing  the  gas  and  to 
prevent  the  acid  from  coming  in  contact  with  skin.  Perform  experiment 
in  the  hood. 

Experiment  III.  —  Formation  of  Silicon  Tetrafluoride  by  the  Action  of 
H2F2  on  Silica. 

Mix  i  gram  of  silica,  SiO2  (sand),  with  i  gram  of  CaF2;  place  the 
mixture  in  a  test  tube;  add  a  little  concentrated  H2SO4;  heat  gently. 
Expose  a  drop  or  film  of  H2O  to  the  action  of  the  fumes  of  Si  F4;  observe 
the  milky  appearance  of  the  water  due  to  the  formation  of  silicic  acid, 
H4Si04. 

(i)  CaF2  +  H2S04  - 


(2)  SiO2    +  2H2F2  —  SiF4          +  2H2O, 

(3)  3SiF4  +  4H20   ->  2H2Si  F6  +  H4Si  O4. 

The    fluorides    of    lithium,    sodium,    potassium,    ammonium,    silver, 
mercury,  iron,  aluminum  and  tin  are  soluble  in  water. 

CHLORINE    Cl. 

At.  Wt.  35.45     Mol.  Wt.  70.9. 
Preparation  and  Properties  oj  Chlorine. 

Experiment  I.  —  Review  the  previous  experiments  with  chlorine.    Fix 
in  mind  the  physical  and  chemical  properties  of  the  elementary  substance. 
Hydrogen  Derivatives. 


154  EXPERIMENTAL  CHEMISTRY. 

Experiment  II. — Preparation  and  Properties  of  Hydrogen  Chloride. 

Recall  the  facts  about  hydrogen  chloride.  Compare  its  properties 
with  the  corresponding  hydrogen  compounds  of  the  other  halogens. 
These  hydrogen  compounds  are  styled  the  hydrogen  halides. 

H  +  Cl  — HC1  +  22,000  cal. 

All  chlorides,  with  the  exception  of  silver,  mercury  (ous),  lead  and  a 
few  basic  chlorides  like  BiOCl  and  SbOCl  are  soluble  in  water.  PbCl2 
is  slightly  soluble  in  cold  water  and  easily  soluble  in  hot  water. 

Oxygen  Derivatives  oj  Chlorine. 

Experiment  III. — Preparation  of  Sodium  Hypochlorite  and  Hypo- 
chlorous  Acid.  Oxidizing  Power  of  Hypochlorous  Acid — Bleaching. 

(a)  Dissolve  3  grams  of  NaOH  in  20  cm3,  of  H2O;  pass  chlorine  into 
the  cold  dilute  solution,  but  do  not  saturate  it.     The  probable  reaction 
may  be  indicated  by  the  following  equation: 

2NaOH  +  C12  -»  Nad  +  NaOCl  +  H2O. 

(b)  Divide  the  solution  into  two  parts;  to  one  part  add  dilute  H2SO4; 
observe   the   odor;  suspend  a  piece  of  cheap  red  calico  and  a  strip  of 
litmus  paper  in  the  acidulated  solution   for  a  day;  then  observe  the 
bleaching  effect  of  hypochlorous  acid,  HC1O    upon  paper  and  calico. 
Write  equations. 

Ascertain  the  bleaching  properties  of  a  non-acidulated  solution  of 
NaOCl  by  repeating  above,  using  the  second  portion  of  the  solution. 
After  removing  paper  and  calico,  add  a  few  drops  of  the  solution  to  an 
indigo,  C16H10N2O2,  solution.  Effects?  Add  a  few  cm.3  of  dilute 
H2SO4  to  solution,  then  add  3  cm.3  of  the  solution  to  the  indigo;  observe 
the  bleaching  effect.  Equations  explaining  activity  of  HC1  are  as 
follows: 

2HC1O  —  2HC1  +  2O  +  18,600  cal. 
C16H10N2O2  +  2(3  —  2C8H5NO2  +  1,800  cal. 

C16H10N2O2  +  2HC1O  ->  2C8H5NO2  +  2HC1  +  20,400  cal. 

Experiment  IV. — Bleaching  Powder  ("Chloride  of  Lime"). 

Place  5  grams  of  chloride  of  line  (Ca(OCl)2  +  CaCl2)  in  a  flask;  add 
20  cm.3  of  H2O  and  shake  thoroughly;  filter.  Using  the  clear  filtrate, 
repeat  Exp.  III.  Do  not  neglect  to  acidulate  solution  or  moisten  objects 
to  be  bleached  with  very  dilute  H2SO4(i  drop  in  25  cm.3  of  H2O). 

Experiment  V. — Preparation  and  Properties  of  Potassium  Chlorate. 

(a)  Dissolve  3  grams  of  KOH  in  10  cm.3  of  H2O;  saturate  the  boiling 
solution  with  chlorine;  cool  the  solution  and  observe  the  formation  of 
crystals  of  potassium  chlorate,  KC1O3.  Devise  a  method  for  proving 


THE    NON-METALS,    OR   ACID-FORMING    ELEMENTS.  155 

that  this  substance  is  identical  with  the  KC1O3  (solid)  found  upon  the 
end  shelves.     Equations  ? 

(b)  KC1O3  as  an  oxidizer.  Mix  a  small  quantity  of  powdered  KC1O3 
with  a  little  powdered  charcoal;  place  the  mixture  upon  the  cover  of  a 
crucible  and  heat  gently.  Results  ?  When  the  cover  has  cooled,  place 
it  in  a  small  beaker  which  contains  25  cm.3  of  H2O;  allow  the  products  of 
the  fusion  to  dissolve;  filter;  add  a  few  drops  of  the  filtrate  to  2  cm.3 
of  AgNO3;  if  a  white  precipitate  forms  which  turns  dark,  on  exposure 
to  light  it  is  AgCl.  Explain  by  use  of  equations. 

Experiment  VI.  —  Chloric  Acid.     Chlorine  Dioxide. 

(a)  Place  a  small  crystal  of  KC1O3  in  the  bottom  of  a  test  tube;  add 
a  few  drops  of  strong  H2SO4;  point  the  mouth  of  the  tube  away  from 
your  face,  then  heat  cautiously;  the  explosive  action  is  due  to  the  presence 
of  a  yellow  gas,  chlorine  dioxide,  C1O2,  which  is  violently  explosive  owing 
to  its  ready  decomposition  into  chlorine  and  oxygen,  with  liberation 
of  much  heat. 

2KC103  +  H2S04   —  2HC103  +  K2S04, 

4HC1O3         -»  2H2O  +  4C1O2  +  2O, 
2HC1O3  +  2O         ->  2  HC1O4; 
or 

2C1O2        —  »  CT2  +  2O2; 


3HC1O3     —  HC1O4  +  2OO    +  H2O, 
again, 


6KC103  +  3H2S04->  2HC104  +  4C1O2  +  3K2SO4  +  2H2O. 
Chloric  acid  —  Structural  formula:     H  —  O  —  Cl  /^  Q  • 

(b)  Oxidizing  action  of  chloric  acid  (C1O2). 

Powder  i  gram  of  KC1O3;  place  the  powder  upon  a  sheet  of  paper 
and  mix  with  it  i  gram,  of  sugar;  put  the  mixture  upon  an  iron  plate; 
allow  a  drop  of  strong  H2SO4  to  drop  from  the  end  of  a  glass  rod  upon 
the  mixture.  Result? 

Note.—  White  gunpowder  is  a  mixture  of  KC1O3,  K4Fe(CN)6  and 
Ci2H22On. 

HC1O3,  Aq—  >  HC1,  Aq  +  30  +  15,300  cal. 

(c)  Explosive  conduct  of  chlorates.     Place  very  small  quantities  of 
KC1O3  and  sulphur  in  a  mortar;  rub  the  mixture  vigorously  with  the 
pestle.     Sharp  explosions  result  from  the  friction.     As  a  rule,  should 

.  chlorates  be  pulverized  with  other  substances  ? 

All  chlorates  are  soluble  in  water. 

Experiment  VII.  —  Perchlorates  and  Perchloric  Acid. 

Perchlorates  may  be  formed  by  suitably  heating  chlorates;  corre- 
sponding chlorides  will  be  formed  simultaneously.  When  either  NaClO3 
or  KC1O3  is  heated  oxygen  is  evolved;  if  the  operation  is  stopped  when 


156  EXPERIMENTAL  CHEMISTRY. 

one-third  of  the  oxygen  has  been  liberated,  the  chloride  and  the  chlorate 
of  the  metal  remain  in  the  fused  mass.  The  more  soluble  chloride  may 
be  dissolved  out  from  the  mass. 

2  KC103  —  KC104 _+  KC1  +  Q; 
KC104-*KC1  +  202. 

Perchloric  acid,  HC1O4,  may  be  formed  by  the  action  of  strong  H2SO4 
on  perchlorates.  Equation? 

HC1O4,  Aq  —  HC1,  Aq  +  lzO~2  +  700  cal. 
Structural  formula  for  HC1O4  ? 

BROMINE,   Br. 

At.  Wt.  79.96      Mol.  Wt.  159.92. 
Preparation  and  Properties  oj  Bromine. 
Experiment  I. — (Hood.) 

Mix  2  grams  of  powdered  MnO2  with  an  equal  weight  of  pulverized 
NaBr  or  KBr.  Introduce  the  mixture  into  the  test  tube  used  when  pre- 
paring oxygen  from  KC1O3  and  MnO2;  add  a  few  cm.3  of  strong  H2SO4; 
heat  gently  and  conduct  the  gas  into  a  50  cm.3  flask  which  is  nearly  im- 
mersed in  cold  water.  Observe  color,  odor  and  relative  density  of  the  gas. 
When  one  or  two  drops  of  the  bromine  gas  have  condensed  in  flask, 
conduct  the  bromine  into  a  test  tube  containing  15  cm.3  or  20  cm.3  of 
H2O.  Compare  the  drop  of  condensed  bromine  vapor  with  the  liquid 
bromine  found  on  side  shelf.  If  the  H2O  in  the  test  tube  was  not  satu- 
rated with  bromine,  add  a  drop  of  the  liquid  bromine  to  the  bromine 
water  (H20  -f  Br.);  shake  thoroughly.  What  is  the  color  of  the  solution 
(bromine  water)  ?  Test  the  action  of  the  solution  on  litmus  paper.  To 
one  portion  of  the  solution  add  2  cm.3  of  ether;  to  a  second  portion  2  cm.3 
of  carbon  disulphide;  and  to  a  third  portion,  2  cm.3  of  chloroform;  shake 
each  tube  vigorously.  Results? 

The  following  partial  equations  may  assist  in  interpreting  the  reactions 
involved  in  preparing  bromine: 

NaBr  +  H2SO4-> Na2HSO4  +  (H  Br). 
MnO2  +  H2SO4— MnSO4  -KH2O  +  (O). 
2HBr  +  O          <±  H2O  +  Br2. 
or 

4HBr  +  MnO2  —  MnBr2  +  2H2O  +  BF2. 
MnBr2  +  H2SO4  —  MnSO4  +  2HBr. 

Note. — If  larger  quantities  of  bromine  are  to  be  prepared  a  retort  and 
a  receiving  flask  may  be  substituted  for  the  above  apparatus. 

Experiment  II. — Substituting  Power  of  Chlorine  as  Compared  with 
Bromine. 

(a)  Dissolve  a  very  small  crystal  of  NaBr  or  KBr  in  15  cm.3  of  water  in 


THE    NON-METALS,    OR    ACID-FORMING    ELEMENTS.  157 

a  test  tube;  add  a  few  cm.3  of  chlorine  water  (side  shelf);  divide  the 
solution  into  two  parts.  Heat  one  portion  gently  and  note  the  change 
in  color.  Explain. 

(b)  To  the  second  portion  add  2  cm. 3  of  CS2;  shake  thoroughly. 
Results?  Equations?  Which  of  the  two — chlorine  or  bromine— has 
the  greater  affinity  for  sodium  ? 

Exp.  III. — Displacement  of  Sulphur  from  Hydrogen  Sulphide. 

Prepare  a  saturated  solution  of  bromine  by  adding  a  drop  of  liquid 
bromine  to  10  cm.  of  H2O;  pass  H2S  through  the  solution  until  the  brown 
color  due  to  the  presence  of  bromine  has  disappeared.  Is  there  any 
evidence  of  the  presence  of  free  sulphur?  Filter  the  solution;  boil  the 
nitrate  until  the  fumes  will  not  darken  a  piece  of  filter  paper  moistened 
with  a  solution  of  lead  nitrate;  test  the  nitrate  with  the  litmus  paper. 
To  2  cm.3  of  AgNO3  add  H2S;  repeat,  using  a  few  drops  of  the  above 
nitrate.  Equations  for  all  reactions?  What  substance  is  produced 
by  the  displacement  of  sulphur  from  H2S  by  bromine?  Recall  the 
similar  conduct  of  chlorine  in  an  analogous  experiment. 

Hydrogen  Derivatives. 

Experiment  IV. — (Hood.)  Preparation  and  Properties  of  Hydrogen 
Bromide — Preliminary. 

Pulverize  about  3  grams  of  NaBr  or  KBr,  and  place  it  in  a  test  tube; 
add  sufficient  strong  H2SO4  to  thoroughly  moisten  powder;  heat  gently; 
blow  your  breath  across  the  mouth  of  the  tube.  Results  ?  Dip  a  glass 
rod  into  NH4OH,  then  hold  rod  near  mouth  of  test  tube.  Results? 
Hold  a  moistened  piece  of  blue  litmus  paper  in  the  evolved  gas.  Results  ? 
What  is  the  color  of  hydrobromic  acid,  HBr? 

What  hydracids  studied  previously  conduct  themselves  in  a  similar 
manner  toward  the  identical  tests?  Write  the  equations  for  the  above 
reactions. 


NaBr  +  H2SO4  —  NaHSO4  +  HBr. 

What  is  the  effect  of  increasing  the  temperature  of  the  mixture  in  the 
test  tube  ?  What  is  the  colored  gas  which  makes  its  appearance  ? 

NaBr  +  H2SO4  —  NaHSO4_+  (HBr), 
2HBr  +  H2SO4  ->  SO2  +  Br2  +  H2O. 

H2SO4  acts  as  an  oxidizing  agent.  Was  either  free  fluorine  or  chlorine 
observed  during  the  preparation  of  their  respective  hydracids  from 
corresponding  salts  by  the  use  of  H2SO4? 

Experiment  V. — Laboratory  Source  of  Hydrogen  Bromide.  Hydro- 
bromic Acid. 

Into  a  test  tube  which  has  been  provided  with  a  stopper  and  a  de- 
livery tube,  introduce  5  grams  of  powdered  NaBr  or  KBr;  add  sufficient 
concentrated  phosphoric  acid,  H3PO4,  to  cover  the  powder;  shake  until 


158 


EXPERIMENTAL  CHEMISTRY. 


the  two  substances  are  mixed  thoroughly;  fit  stopper  and  delivery  tube 
into  place;  heat  tube  gently;  collect  the  evolved  gas  in  5  cm.3  of  H2O 
in  a  test  tube  the  mouth  of  the  delivery  tube  should  be  about  .2  cm. 
above  the  surface  of  the  water. 

Is  there  any  evidence  of  the  presence  of  free  bromine?  In  what 
particular  respect  is  H2SO4  different  from  H3PO4  in  its  action  on  bro- 
mides ?  The  acid  solution  may  be  tested  as  in  Experiment  IV.  Add 
a  few  drops  of  the  solution  to  a  little  sodium  carbonate.  Results? 
Repeat  the  foregoing  experiment,  using  a  few  cm.3  of  AgNO3.  Results  ? 


FIG.  28. 

If  the  acid  solution  is  saturated,  its  effect  on  a  piece  of  zinc  should  be 
ascertained.  Wrap  a  piece  of  platinum  wrire  around  the  zinc.  Why? 
Equations  ? 

H  +  Br  -*  HBr  +  8400  cal. 
Experiment  VI. — Optional.     Preparation  of  Pure  Hydrobromic  Acid. 

Note. — When  phosphorous  and  bromine  are  mixed,  they  combine 
energetically  forming  phosphorous  tribromide  (PBr3).  Pure  HBr  is 
prepared  by  hydrolysis  of  this  substance. 


/Br  — H— OH 


/OH 


P  — Br  —  H  —  OH  —  P— OH  +  3  H  —  Br. 
\Br-H—  OH         \OH 

The  preparation  of  the  acid  may  be  conducted  as  follows :  The  ordinary 
funnel  of  a  300  cm.3  generating  flask  is  displaced  by  a  dropping  funnel;  the 
flask  is  connected  in  series  with  two  U- tubes  (or  see  Fig.  28) ;  the  U-tube  next 
to  the  flask  should  be  half  filled  with  glass  beads  which  have  been  moist- 
ened and  rolled  in  red  phosphorous;  the  second  tube  should  be  about 
one-fourth  filled  with  water  to  absorb  the  HBr.  The  mouth  of  the  tube 


THE    NON-METALS,    OR   ACID-FORMING    ELEMENTS.  159 

delivering  the  gas  to  the  tube  should  be  above  the  surface  of  the  water. 
Charge  the  generating  flask  with  5  grams  of  red  phosphorus  and  enough 
water  to  barely  cover  the  latter.  (It  is  suggested  that  5  grams  to  10  grams 
of  clean  sand  be  added  with  the  charge.)  Pour  10  cm.3  of  bromine 
(liquid)  into  the  funnel;  when  apparatus  has  been  assembled  properly, 
allow  the  bromine  to  fall  on  the  phosphorus  drop  by  drop.  The  union 
of  the  bromine  and  phosphorus  will  be  accompanied  by  a  flash  of  light. 
(Can  you  suggest  a  reason  for  using  red  instead  of  yellow  phosphorus  ?) 
After  the  water  in  the  flask  has  been  saturated  with  the  HBr,  the  latter 
will  escape  into  the  U-tube  and  eventually  be  absorbed  in  the  water  con- 
tained in  the  most  remote  U-tube.  If  too  little  water  is  present  in  the  flask, 
crystals  of  PBr3  will  clog  the  apparatus;  the  addition  of  a  few  drops  of 
water  will  prevent  their  formation.  The  acid  properties  of  the  solution 
of  the  gas  may  be  tested  as  previously.  If  convenient  its  conductivity 
may  be  -determined. 

The  bromides  are  easily  soluble'  in  water  with  the  exception  of  silver, 
mercury  (ous)  and  lead.  PbBr2  is  slightly  soluble  in  cold  water  and 
easily  soluble  in  hot  water. 

Oxygen  Derivatives  oj  Bromine. 

Note. — No  oxides  of  bromine  have  been  prepared. 

Experiment  VII. — Preparation  of  Sodium  Hypobromite  and  Hypo- 
bromous  Acid. 

(a)  Dissolve  3  grams  of  NaOH  in  20  cm.3  of  H2O;  add  a  few  drops  of 
liquid  bromine  to  the  cold  dilute  solution.     Observe  the  disappearance 
of  the  color  of  the  bromine.     The  bromine  should  be  added  drop  by 
drop,  but  not  in  a  sufficient  quantity  to  saturate  the  solution. 

2NaOH  +  Br2  —  NaBr  +  NaBrO  +  H2O. 

(b)  Divide  the  above  solution  into  two  parts.     To  one  portion  add 
dilute  H2SO4;   observe  the  color  and  odor  of  the  evolved  gas.     The 
following  partial  equations  probably  represent  the  stages  of  the  reaction: 

NaBr  +  NaBrO  +  H2SO4 -*  Na2SO4  +  (HBr)  +  (HBrO). 
HBr    +  HBrO    —  H2O  +  Br2. 
Adding,  NaBr  +  NaBrO  +  H2SO4  ->  Na^SO,  +  H2O  +  Br2. 

(c)  Repeat  (b)  using  acetic  acid  instead  of  H2SO4.     Use  slight  excess 
of  acid.     Note  that  bromine  is  liberated.     Write  the  structural  formula? 
for  NaBrO  and  NaBrO3. 

Experiment  VIII. — Preparation  and  Properties  of  Potassium  Bromate. 

Dissolve  3  grams  of  KOH  in  10  cm.3  of  H2O;  saturate  the  solution 
with  bromine;  heat  to  boiling.  When  crystals  make  their  appearance, 
dissolve  them  in  H2O;  add  a  few  drops  of  liquid  bromine;  boil  to  crystal- 
lization; dissolve  crystals  in  the  smallest  quantity  of  water;  pour  the 


l6o  EXPERIMENTAL  CHEMISTRY. 

solution  into  a  crystallizing  dish.     The  KBrO3  crystals  may  be  separated 
from  the  more  soluble  KBr  crystals  by  crystallization.     Equation? 

Experiment  IX. — Bromic  Acid.  Interaction  of  Bromic  and  Hydro- 
bromic  Acids. 

(a)  To  a  solution  of  NaBrO3  or  KBrO3,  add  dilute  H2SO4.     Is  bro- 
mine liberated?     Equation? 

(b)  To  a  solution  containing  KBrO3  add  dilute  H2SO4.     Is  bromine 
liberated  ?     Equations  ? 

The  following  thermo-chemical  equation  gives  an  approximate  idea 
of  the  oxidizing  power  of  bromic  acid: 

HBrO3,  Aq— *  HBr,  Aq  +  30  +  15,000  cal. 
Structural  formula  for  hydrobromic  acid: 

O 

S 

H— O  — Br 

\ 
O 

IODINE.     I. 

At.  Wt.  126.97         Mol.  Wt.  254.8  (i85°-6oo°). 
Preparation  and  Properties  oj  Iodine. 

Experiment  I.— 

(a)  Mix  2  grams  of  powdered  MnO2  with  an  equal  weight  of  pulverized 
Nal  or  KI;  place  the  mixture  in  an  evaporating  dish;  add  sufficient 
concentrated  H2SO4  to  moisten  the  mixture  thoroughly;  clamp  an  in- 
verted funnel  over  the  evaporating  dish  in  such  a  position  that  all  vapors 
from  the  dish  will  escape  through  the  stem  of  the  funnel.     Heat  the 
dish  gently.     A  dense  vapor  will  be  set  free.     Observe  color,  odor  and 
relative  density  of  the  vapor  before  it  condenses  upon  the  sides  of  the 
funnel.     Continue  to  heat  dish  until  colored  vapor  ceases  to  be  evolved. 
After  cooling  funnel,  scrape  crystals  into  a  clean  beaker   to  be  used  for 
following  experiments. 

(b)  Resublimed  iodine.     Place  a  few  of  the  crystals  in  a  flask;  heat 
the  bottom  of  the  flask  gently  to  volatilize  the  iodine  which  crystallizes 
upon  the  cooler  portions  of  the  flask.     Does  the  iodine  fuse  before  be- 
coming a  vapor?     What  are  such  processes  called? 

(c)  Test  the  solubility  of  iodine  in  each  of  the   following   solvents: 
H2O,   C2H5OH,   CHC13,  (C2H5)2O,   CS2,  and  a  KI  solution.     Recall   a 
similar  experiment  under  subject  of  "  Solutions." 

(d}  Triturate  .3  gram  of  starch  and  20  cm.3  of  H2O  in  a  mortar  until 
they  are  thoroughly  mixed.  Pour  this  mixture  into  100  cm.3  of  boiling 
water;  boil  the  mixture  for  about  five  minutes.  Add  a  few  drops  of  the 
cold  starch  emulsion  to  a  few  cm.3  of  a  KI  solution  of  iodine;  heat  the 
mixture.  Does  the  color  disappear  ?  Cool.  Results  ?  The  compound 


THE    NON-METALS,    OR    ACID-FORMING    ELEMENTS.  l6l 

formed  with  the  characteristic  blue  color  is  known  as  starch  iodide. 
Iodine  is  used  to  detect  the  presence  of  starch^  and  vice  versa. 

Note. — Bottle  the  starch  emulsion  and  preserve  it  for  future  use. 

The  solution  of  iodine  in  alcohol  is  known  as  the  tincture  of  iodine. 
The  brown  color  of  the  tincture  is  attributed  to  the  fact  that  the  iodine 
and  alcohol  are  in  a  loose  state  of  combination;  likewise  iodine  dis- 
solved in  a  KI  solution  is  supposed  to  form  a  definite  compound: 

2KI  +  I2~*  2KI3. 

On  the  other  hand,  the  blue  substance,  starch  iodide,  referred  to  as  a 
compound,  is  not  a^  chemical  compound  but  merely  a  solution  of  iodine 
in  starch. 

Experiment  II. — Substituting  Power  of  Chlorine  and  Bromine  as 
Compared  with  Iodine. 

(a)  Dissolve  a  very  small  crystal  of  Nal  or  KI  in  15  cm.3  of  water  in 
a  test  tube;  divide  it  into  two  portions;  to  one  portion  add  a  few  drops  of 
chlorine  water;  heat  gently  and  observe  the  change  in  color;  add  a  few 
drops  of  the  solution  to  a  little  of  the  starch  solution.     Results  ?     Shake 
up  the  rest  of  the  solution  with  3  cm.3  of  CS2.     Results?     Equations? 

Note. — An  excess  of  chlorine  water  decolorizes  a  solution  of  iodine 
in  CS2. 

(b)  Repeat  (a)  using  the  second  portion  of  the  solution  and  bromine 
water  instead  of  chlorine  water.     Results  ?     Equations  ? 

In  view  of  the  above  and  previous  experiments  what  do  you  infer  in 
regard  to  the  relative  affinities  of  the  halogens  for  hydrogen?  Arrange 
them  in  the  order  of  their  increasing  affinity. 

Experiment  III. — Potassium  Bichromate  Liberates  Iodine  from  an 
Acidified  Solution  of  an  Iodide. 

To  a  dilute  solution  of  KI  add  CS2;  shake.  Results?  Acidify  the 
solution  with  a  little  dilute  H2SO4,  and  add  a  few  drops  of  a  K2Cr2O7* 
solution;  agitate  the  mixture.  Results? 

Note, — Bromine  is  not  liberated  from  bromides  by  above  process. 
Test.  Equations  ? 

Experiment  IV. — Iodine  is  Liberated  by  the  Oxidizing  Action  of  Sul- 
phuric Acid. 

Place  a  small  crystal  of  Nal  or  KI  in  a  test  tube;  add  3  cm.3  of  concen- 
trated H2SO4;  heat  gently.  Observe  the  color  and  odor  of  the  evolved 
gas.  Identify  it.  Ascertain  whether  hydriodic  acid,  HI,  is  evolved  or 
not. 

KI  +  H2S04<=±  KHS04  +  (HI), 
2HI  +  H2S04^±  S02  +  I2  +  H20, 
6HI  +  H2S04<=±  S  +  3I2  +  4H20, 
SHI  +  H2SO4  *±  H2S  +  4l2  +  4H2O. 
*  Sodium  nitrite  may  be  sustituted  for  K2Cr3O7. 
ii 


1 62  EXPERIMENTAL  CHEMISTRY. 

Compare  the  above  reaction  with  the  action  of  H2SO4  on  CaF2,  NaCl, 
NaBr.  Arrange  the  hyrdascids  of  the  halogens  in  the  order  of  their 
increasing  stability  toward  oxidizing  agents. 

Experiment  V. — Displacement  of  Sulphur  from  Hydrogen  Sulphide. 

To  2  grams  of  powdered  iodine  in  a  flask  fitted  with  a  cork  containing 
a  single  perforation  add  20  cm.3  of  H2O;  pass  H2S  through  the  mixture 
until  iodine  disappears.  Proceed  with  experiment  as  per  directions 
Exp.  Ill,  "Bromine." 

Hydrogen  Derivatives. 

Experiment  VI. — Preparation  and  Properties  of  Hydrogen  Iodide- 
Hydriodic  Acid. 

(a)  Follow  directions  given  in  Exp.  V.,  "Bromine,"  substituting  the 
words  iodide  and  iodine  for  bromide  and  bromine,  respectively.     Equa- 
tions ? 

(b)  Pure  HI  may  be  prepared  by  a  process  similar  to  that  used  in  the 
preparation  of  pure  HBr — Exp.  VI.,  "Bromine."     Use  the  same  appa- 
ratus.    The  flask  may  be  charged  with  a  mixture  of  iodine  and  water 
(4  to  i),  and  the  red  phosphorus  (stirred  to  a  paste  with  water)  allowed 
to  fall  slowly  on  the  iodine,  drop  by  drop. 

Another  method,  using  same  apparatus,  is  as  follows:   Mix  the  red 
phosphorus  and  iodine  (10  to  i  by  weight)  in  a  dry  mortar;  charge  flask 
with  this  mixture,  and  allow  water  to  fall  slowly  upon  it  drop  by  drop. 
H  -f  I  — »  HI  —  6000  cals. 

Experiment  VII. — (Optional.)  Use  of  Potassium  Iodide  to  Determine 
Rate  of  Absorption  from  the  Stomach. 

Test  the  saliva  for  iodine  by  the  iodide  of  starch  test.  Take,  by  mouth, 
a  capsule  containing  0.2  gram  of  KI  or  10  drops  of  the  saturated  solution. 
Note  the  time.  Test  the  saliva  at  end  of  each  minute  for  iodine.  If 
absorption  is  normal,  a  positive  test  should  appear  in  fifteen  minutes. 
Taking  equal  quantities  of  the  iodide,  the  members  of  the  class  will  be 
able  to  make  a  comparison  of  the  rates  of  absorption.  Determine  the 
average  rate  of  absorption  of  the  individuals  of  the  class  who  are  fleshy; 
likewise  calculate  the  average  rate  of  those  who  are  not  fleshy.  Your 
conclusions  ? 

DERIVATIVES    OF    IODINE. 

Oxygen. 

Note. — Although  iodine  and  ordinary  oxygen  do  not  combine  with 
one  another,  O3  reacts  with  iodine  to  form  I2O5.  Compare  with  con- 
duct of  fluorine. 

Experiment  VIII. — Preparation  of  Sodium  Hypoiodite.    Sodium  lodate. 

These  salts  may  be  prepared  by  a  method  analogous  to  that  used  in 
the  preparation  of  the  corresponding  salts  of  bromine  and  chlorine. 


THE    NON-METALS,    OR   ACID-FORMING    ELEMENTS.  163 

The  hypoiodite,  however,  oxidizes  quite  readily,  forming  the  iodate  and 
iodide.  If  this  mixture  is  treated  with  acid,  iodic  and  hydriodic  acids 
are  liberated,  but  interact,  liberating  iodine.  Recall  the  conduct  of 
corresponding  acids  of  bromine.  The  oxacids  of  iodine,  HIO  and  HIO3, 
are  quite  stable.  Hydriodic  acid: 

O 

^ 
H  — O  — I 

\ 
O. 

Experiment  IX. — Relative  Affinity  of  Bromine  and  Iodine  for  Oxygen. 

To  5  cm.3  of  a  solution  of  KBrO3  add  5  cm.3  of  dilute  H2SO4. 
Divide  the  solution  into  two  parts.  To  one  portion  add  CS2  and  shake. 
Is  the  CS2  colored?  To  the  other  portion  add  a  very  small  flake  of 
iodine;  shake  vigorously  for  two  or  three  minutes;  decant  the  clear 
solution  into  another  test  tube;  add  a  few  drops  of  CS2;  agitate  the  con- 
tents. Does  the  reddish-brown  color  indicate  that  iodine  has  a  greater 
affinity  than  bromine  for  oxygen  ?  Your  reasons  ?  (Examine  structural 
formula  of  HBrO3.) 

Experiment  X. — Properties  of  lodates. 

(a)  Heat  a  few  crystals  of  KIO3  or  NaIO3  in  a  test  tube  and  test  for 
oxygen.     Recall  the  preparation  of  oxygen  from  KC1O3.     Equation? 

(b)  Pulverize  .5  gram  of  KIO3  or  NaIO3;  mix  thoroughly  with  .3  gram 
of  powdered  charcoal;  place  the  mixture  in  a  crucible  and  heat  until  the 
entire  mass  glows;  cool;  place  crucible  and  contents  in  a  beaker  of  hot 
water  to  dissolve  the  fused  mass.     Filter  the  solution.     Devise  a  method 
to  prove  that  an  iodide  was  formed  by  the  above  reaction. 

2NaIO3  +  3C  —  2NaI  +  3CXX. 

(c)  Place  a  small  crystal  of  an  iodate  in  a  mortar  with  a  little  sulphur; 
rub  the  mixture  vigorously  with  a  pestle.     Sharp  explosions  result  from 
the  friction.     What  other  substance  reacts  like  the  iodate  ?     Equations  ? 
Would  you  infer  that  iodates  are  strong  oxidizers?     Structural  formula 
of  NaIO3  ? 

Experiment  XI. — Preparation  of  Iodic  Acid. 

Note. — Iodic  acid  may  be  prepared  by  a  method  analogous  to  the 
preparation  of  chloric  acid,  i.e.,  by  the  action  of  H2SO4  on  a  soluble 
iodate. 

2NaIO3  +  H2SO4  —  Na2SO4  +  2HIO3. 

The  iodic  acid  may  be  crystallized  out  from  the  solution  as  a  white 
solid. 

The  author  has  found  the  following  method  for  the  preparation  of 
the  acid  to  be  very  satisfactory:  Place  two  or  three  small  flakes  of 


164  EXPERIMENTAL  CHEMISTRY. 

iodine  in  the  bottom  of  the  largest  test  tube'among  your  apparata;  add 
15  cm.3  of  strong  HNO3;  take  the  mixture  to  the  hood  and  boil  vigorously 
until  the  solution  becomes  colorless.  It  may  be  necessary  to  add  HNO3 
from  time  to  time  to  oxidize  all  of  the  iodine.  The  colorless  solution  of 
HIO3  should  be  boiled  for  sometime  to  concentrate  it,  and  then  set  aside 
to  allow  crystals  of  the  acid  to  form,  or  the  solution  may  be  tested  at 
once  as  follows: 

Add  a  few  drops  of  HNO3  to  an  AgNO3  solution.  Does  a  precipitate 
form  ?  Repeat  foregoing  using  a  solution  of  KI  instead  of  the  HNO3  ;  ob- 
serve the  yellow  color  of  the  precipitate,  Agl.  Now  add  a  few  drops  of 
the  prepared  solution  supposed  to  contain  iodic  acid,  to  3  cm.3  of  the 
AgNO3  solution.  What  is  the  color  of  the  precipitate  ?  Was  it  formed 
by  the  action  of  HNO3  ?  Your  reasons  ? 


3-*  2H2O  +  4NO2  +  (O2). 
I2  +  H2O  +  50  ->  2HIO3. 

Save  the  iodic  acid  solution  for  the  following  experiment. 
Experiment  XII.  —  Oxidizing  Action  of  Iodic  Acid. 

To  5  cm.3  of  the  solution  of  iodic  acid  add  2  cm.3  of  CS2;  mix  thoroughly. 
Note  that  the  color  of  the  CS2  remains  unchanged.  Add  a  few  drops 
of  sulphurous  acid,  H2SO3  to  the  solution  and  then  shake  thoroughly. 
What  is  the  substance  which  imparts  the  color  to  the  CS2?  Has  the 
odor  of  SO2  disappeared? 

2HIO3  +  5H2SO3->  sH2SO4  +  I2  +  H2O. 
HIO3,  Aq  —  »  HI,  Aq  +  36  —  42,900  cal. 

Experiment  XIII.  —  (L.  T.)  Periods  of  Induction  in  Chemical  Reac- 
tions. 

Instructions.  —  "There  is  a  pleasing  'lecture  experiment'  for  illus- 
trating the  'period  of  induction.'  A  very  dilute  solution  of  H2SO3 
and  HIO3  (i  gram,  e.g.,  in  600  liters  of  water)  is  mixed  with  starch.  The 
appearance  of  a  visible  blue  color  occupies  a  measurable  time,  which 
may  be  extended  by  using  more  dilute  solutions."  —  Mellor's  "Chemical 
Statics  and  Dynamics." 

Experiment  XIV.  —  Resemblance  of  the  Properties  of  Corresponding 
Compounds  of  the  Halogens. 

(a)  To  a  solution  of  Pb(NO3)2  add  a  few  drops  of  a  KC1  solution; 
observe  color  of  precipitate;  filter.     Determine  solubility  of  precipitate 
in  hot  water.     Repeat  above  using  KBr.     Also  KI. 

(b)  Repeat  (a)  substituting  AgNO3  for  Pb(NO3)2,  and  NH4OH  for 
hot  water. 

(c)  Repeat  (b)  substituting  HgNO3  for  AgNO3. 

Note  the  general  similarity  of  conduct  of  the  respective  halides  of  each 
metal.  Tabulate  results. 


THE    NON-METALS,    OR    ACID-FORMING    ELEMENTS. 


Experiment  XV. — Detection  of  Fluorides,  Chlorides,  Bromides  and 
Iodides  together  in  a  Solution. 

Devise  a  system  of  tests  which  will  provide  for  the  detection  of  the 
halogens  in  solution.  Apply  the  system  to  a  solution  containing  the 
four  halides.  Ask  the  assistant  to  give  you  an  " unknown"  solution; 
test  for  the  presence  of  the  halides.  Report  (  ?).  Make  a  complete  record 
of  all  work. 

A  general  outline  of  the  relations  of  the  halogens  is  presented  in  the 
following  table: 

Physical  Properties  Fluorine  Chlorine  Bromine  Iodine 

Atomic  wreight  19  35-45                79-9^  126.85 

State  or  phase  Gas  Gas  Liquid  Solid 
Color 


Specific  gravity 
Specific  heat 
Molecule  of  gas 
Melting  point 
Boiling  point 
Chemical  Properties 
Water  decomposed 


Fluorine 

Chlorine         Brc 

Gas 

35-45                7^ 
Gas               Lii 

Pale 

Greenish- 

yellow 
15  (liquid) 

yellow            Brc 
1.15  (liquid)        3. 

Black 


F2  C12 

—102°  (solid) 


—  186° 


-70     >7 

—  33-7 


Br2    I2(2oo°-6oo°) 

o 
l84° 


7-3°          H4° 


59 


In  the  dark,    In  the  sun-     Slowly  in  Only  by  in- 
light,            the  sun-  direct  pro- 
light,  cesses. 

Readily,      In  the  sun-     In  heat  of  Only  by  in- 
light,              flame,  direct  pro- 
cesses. 

H  F             HC1                H  Br  HI 

37,600  cal.    22,000  cal.       8400  cal.  — 6000  cal 


Formation  of  H- 
derivatives 

Heat  of  formation 
of  H-derivatives 

Stability  of  H- 
derivatives 
Affinity  for  oxygen 
Oxides 

Oxacids 

V 

I* 
V 

*  The  stability  of  the  oxacids  and  their  corresponding  salts  increases  with  increase 
jn  oxygen-content. 


^v 

X, 

v 

s 

None 
None, 

C120,C1203, 
C1O2. 
HC10 
HC1O2 
HC1O3 
HC104 

IA 

HTCr  O 

H  Br  O3     H 
H  Br  O4  ?  H 

10, 
10. 

1 66  EXPERIMENTAL  CHEMISTRY. 

It  must  be  obvious  to  the  student  that  the  elements  of  this  group  or 
family  are  connected  by  properties,  physical  and  chemical,  which  show 
a  remarkable  similarity.  As  a  matter  of  fact,  the  uniformity  is  so  marked 
that  if  the  elements  are  arranged  so  as  to  show  a  gradation  of  any  one 
property,  the  order  will  not  be  altered  if  they  are  arranged  with  respect 
to  other  properties.  Again,  these  properties  seem  to  bear  so  definite  a 
relation  "to  the  atomic  weight  of  an  element,  that  it  is  now  generally 
believed  by  chemists,  although  it  is  not  yet  fully  proved,  that  the  proper- 
ties of  the  elements,  as  well  as  the  properties  of  their  compounds,  are 
periodic  junctions  oj  their  atomic  weights.  For  example,  the  specific  heat 
of  an  element  is  inversely  proportional  to  its  atomic  weight.  These  and 
other  facts  have  led  to  the  grouping  of  the  elements  into  "  natural  families  " 
or  groups.  The  halogens  are  unique  in  the  fact  only  that  they  represent 
the  most  prominent  example  of  the  gradation  in  properties  observed 
among  the  members  of  a  natural  group  of  elements. 

The  regular,  periodic  variation  in  the  properties  of  the  elements  con- 
stitutes what  is  known  as  the  periodic  law,  by  means  of  which  the  elements 
are  given  their  natural  classification.  The  law  is  illustrated  by  the 
following  tables.  (See  Classification  of  the  Elements.) 

It  should  be  noted  before  leaving  this  group  that  the  halogen  elements 
are,  par  excellence,  acid  elements;  not  only  do  their  oxygen  compounds, 
but  their  hydrogen  compounds  also  exhibit  acid  properties. 


CHAPTER  XVI. 
CLASSIFICATION  OF  THE  ELEMENTS. 

The  classification  of  the  elements  into  metals,  non-metals  and  metal- 
loids was  made  ^during  the  early  development  of  chemistry  when  com- 
paratively few  elements  were  known.  This  classification  was  convenient, 
but  it  soon  led  to  confusion,  because  elements  were  discovered  which 
did  not  fall  exclusively  into  any  one  of  these  three  groups.  It  soon 
became  evident  to  the  chemists  of  that  time  that  a  new  plan  of  classi- 
fication was  necessary;  especially  so  after  the  announcement  of  the 
"Law  of  Multiple  Proportion"  by  Dalton.  As  the  result  of  the  accept- 
ance of  this  law,  Prout  (1815)  promulgated  the  theory  that  the  atomic 
weights  of  all  the  elements  are  exact  multiples  of  the  atomic  weight  of 
hydrogen,  assuming  that  the  atoms  of  other  elements  are  merely  con- 
densations or  aggregations  of  hydrogen  atoms.  The  atomic  weight 
determinations  of  Berzelius,  Marignac  and  Stas  did  not  prove  that  the 
atomic  weights  of  the  elements  are  exact  multiples  of  the  atomic  weight 
of  hydrogen. 

Note. — "  The  work  of  T.  W.  Richards  on  the  atomic  weights  of  a 
large  number  of  the  metals  should  receive  special  attention.  He  has 
improved  old  methods,  devised  new  ones  and  applied  them  with  a  skill 
which  is  rare.  His  determinations  are  to  be  ranked  among  the  very 
best  which  have  ever  been  made." — Jones'  "Physical  Chemistry." 

The  theory  of  Prout,  though  proven  untenable,  was  instrumental  in 
directing  a  great  deal  of  attention  to  the  question  of  atomic  weights,  with 
the  result,  that  Dobereiner  (1825)  "on  examining  the  atomic  weights 
of  correlated  elements,  observed  that  the  atomic  weight  of  the  middle 
member  of  a  group  of  three  related  elements  was  nearly  a  mean  of  the 
atomic  weights  of  the  other  two  elements.  These  groups  of  three  were 
known  as  the  "Triads  of  Dobereiner"  and  probably  represent  the  first 
attempt  to  classify  the  elements  on  the  basis  of  a  relationship  existing 
between  properties  and  atomic  weights.  The  student  is  referred  to 
Venable's  "Development  of  the  Periodic  Law." 

NON-METALS. 

Helium        Fluorine         Oxygen  Nitrogen  Carbon 

Neon  Chlorine        Sulphur  Phosphorus  Silicon 

Argon          Bromine        Selenium         Arsenic  Germanium* 

Krypton       Iodine  Tellurium       Antimony*  Tin* 

Xenon  Bismuth*  Boron 

(Aluminium) 
*  Metalloids. 

167 


i68 


EXPERIMENTAL  CHEMISTRY. 


METALS. 

See  table  of  elements  for  complete  list  of  metals,  non-metals  and  metal- 
loids. 

TRIADS    OF    DOBEREINER    (1825). 

(Examples). 


Lithium 

Sodium 

7 

23 

Potassium 

Rubidium 

39 

8S 

Calcium 

Strontium 

40 

88 

Sulphur 
32 


Selenium 

77 


Potassium 
39 

Caesium 
133 

B  arium 

137 

Tellurium 
127 


7  +  39 


=  23 


39  +  133 


40  +  137 


86 


32  +  127 


=  78 


Chlorine 

35 


Bromine 

80 


Iodine 
126 


35 


=  80 


H. 
Li. 


1  F. 

2  Na. 


8  Cl. 


9  K.     16  Cu.  23  Rb. 
G.    3  Mg.  10  Ca.    17  Zn.   24  Sr 
Bo.  4  Al. 
C.    5  -Si. 
N.    6  P. 
O.    7  S. 


(Newland's,  1864.) 

Co.   & 
15  Ni.  22  Br.  29  Pd.  36  I. 

30  Ag.  37  Cl.         ^   _.    J0 
3iCd.38Ba.  &V.45  Pb.   54 

_     TT  T»  .  £       T'l-          ^  £ 


Pt.   & 

42  Ir.    50 
44  Tl.    53 

10  ^a.     17  z,n.    24  sr.  31  (_,a.  30  &&.  &V-45  Pb. 

11  Cr.    18  Y.     25  Ce.  &La.  32  U.    40  Ta.        46  Th.  56 

12  Ti.     19  In.    26  Zr.  33  Sn.  39  W.         47  Hg.  52 

13  Mn.  20  As.    27  Di.  &Mo.  34  Sb.  41  Nb.        48  Bi. 

14  Fe.    21  Se.    28  Ro.  &  Ru.  35  Te.  43  Au. 


49  Os. 


55 
51 


CLASSIFICATION    OF    THE    ELEMENTS. 


169 


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CLASSIFICATION    OF    THE    ELEMENTS. 


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CLASSIFICATION    OF    THE    ELEMENTS. 


173 


Many  attempts  have  been  made  to  so  modify  the  form  of  the  fore- 
going tables  that  the  position  of  an  element  will  more  faithfully  represent 
its  relationships.  The  following  is  an  example: 


BAYLEY  S    TABLE. 

The  following  tabulated  data  may  serve  to  make  more  significant 
the  relative  positions  of  the  elements  in  the  periodic  table: 

Specific  Gravity. 
Na    Mg    Al     Si    P     S     Cl 

.97         I.y         2.5     2.4    2.2    2.0     1.3 

K.  Ca  Sc  Ti  V.  Cr  Mn  Fe  Co  Ni  Cu  Zn  Ga  Ge  As  Se  Br 

.87  1.6  3.8  4.  ?  5.5  6.8   7.2  7.9  8.5  8.9   8.9  7.2    5.9  5.5  5.6  4.8  3.2. 


174 


EXPERIMENTAL  CHEMISTRY. 


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3  £  £  £ 

CLASSIFICATION    OF    THE    ELEMENTS. 

ATOMIC  VOLUME— SPECIFIC  GRAVITY. 


Element. 

ii 

0    <u 

<£ 

Specific 
Gravity.* 

Atomic 
Volume.* 

Element. 

'g  bo 
o'u 

<^ 

Specific 
Gravity. 

Atomic 
Volume. 

Element. 

>%£ 

S."> 

B£ 

<& 

o  ^ 

<JH     •£ 
11 

£5 

•|1 

S'o 
<> 

Li 

7.0 

°-59 

11.9 

S 

32.0 

2  .05 

!6. 

Cu 

63.6 

8.9 

7-i3 

Be 

9.0 

2.07 

4-4 

Cl 

35-4 

Gas 

26. 

Zn 

65-4 

7-2 

9-37 

B 

II  .0 

2-5 

4.1 

K 

39-° 

0.87 

45-4 

Ga 

70.0 

5-9 

u-5 

C 

12  .0 

3-5 

3-4 

Ca 

40.0 

1.58 

25-3 

As 

75- 

5-6 

13-1 

N 

14.0 

Gas 

Sc 

44-0 

3-8? 

Se 

79.0 

4.8 

16.5 

0 

16.0 

Gas 

Ti 

48.0 

4-? 

12.5? 

Br 

79-9 

3-K; 

25- 

F 

19.0 

? 

V 

51.2 

5-5 

9-3 

Rb 

85-5 

1.52 

56-3 

Na 

23.0 

0.97 

23-7 

Cr 

52  .0 

6.8 

7-65 

Sr 

87.6 

2-5 

34-4 

Mg 

24.0 

1.74 

13.8 

Mn 

5-50 

7.14 

7.6 

Ag 

107.9 

10.6 

10.2 

Al 

27.0 

2.60 

10.6 

Fe 

55-9 

7.86 

6-9 

I 

126.9 

4-95 

25.6 

Si 

28.0 

2-39 

10.7 

Ni 

58.7 

8.90 

6-31 

Te 

127  .6 

6.25 

20-4 

P 

31.0 

2  .  20 

12.8 

Co 

59-o 

8-5 

6.82 

Cs 

132.9 

1.9 

70. 

*Huth's,  Das  periodische  Gesetz  der  Atomgewichte. 


EXERCISES. 

1.  Draw  a  periodic  curve  of  the  elements  named  in  the  above  table, 
tabulating  the  specific  gravities  on  the  vertical  axis,   and   the  atomic 
weights  on  the  horizontal  axis. 

2.  Repeat  (i),  substituting  atomic  volumes  for  specific  gravities. 

3.  Repeat  (i),  substituting  specific  heats  (see  Appendix)  for  specific 
gravities. 


CHAPTER  XVII. 
OXYGEN  FAMILY. 

Oxygen,  O,  16.00, 

Sulphur,  S,  32.06, 

Selenium,  Se,  79.2, 

Tellurium,  Te,  127.6. 

It  has  been  pointed  out  in  the  preceding  chapter  that  the  chemical 
and  physical  properties  of  the  elements  bear  a  definite  relation  to  their 
atomic  weights.  The  student  should  note  carefully  the  conduct  of  these 
four  elements  whose  atomic  weights  place  them  in  the  same  "natural 
group,"  to  ascertain  whether  they  conform  to  this  so-called  "periodic 
law." 

OXYGEN,    O. 

At.  Wt.  16.     Mol.  Wt.  32. 
Preparation  and  Properties  of  Oxygen. 

Experiment  I. — Recall  the  previous  experiments  with  oxygen.  Fix  in 
mind  the  physical  and  chemical  properties  of  the  elementary  substance. 

Hydrogen  Derivatives. 

Experiment  II. — Recall  experiments  suggested  by  the  following  equa- 
tions: 

H2  +  O  — >  HO,   Aq  +  68,360  cals. 
H2  +  °2  — *  H2°2>  Aq  +  45.300  cals. 

Experiment  III. — Oxidizing  Action  of  Hydrogen  Dioxide. 

Place  5  cm. 3  of  lead  acetate  in  an  evaporating  dish;  moisten  a  strip 
of  writing  paper  with  the  solution,  then  hold  the  paper  in  a  current  of 
hydrogen  sulphide.  Results?  Allow  the  paper  to  dry.  Place  2  cm.3 
of  H2O2  in  a  test  tube;  using  a  glass  rod  wet  with  the  H^Oj,  write  your 
name  upon  the  paper.  Explain. 

Oxygen  Derivatives. 

Experiment  IV. — Recall  the  formation  of  ozone. 

O2  +  O  — »  O3  —  32,400  cals. 


OXYGEN    FAMILY.  177 

SULPHUR,    S. 

At.  Wt.  32.06     Mol.  Wt,  (S2>  S4,  Se,  S8). 
Properties  oj  Sulphur. 
Experiment  I. — Transition  Points.     Plastic  Sulphur. 

Procure  a  few  grams  of  roll  sulphur;  note  its  color,  taste,  odor,  relative 
density,  structure,  luster.  Place  5-10  grams  of  the  sulphur  in  a  dry  test 
tube;  heat  gently,  observing  the  yellow  deposit  of  sulphur  (flowers  of  sul- 
phur) on  the  cooler  portions  of  the  tube  as  the  experiment  proceeds ;  increase 
the  temperature  of  the  tube  until  the  sulphur  melts.  Determine  the 
melting  point  by  means  of  the  thermometer,  then  remove  it;  continue 
to  increase  the  supply  of  heat,  noting  the  change  in  color  of  the  fluid. 
At  1 60  °- 2 60°  the  yellow  mobile  fluid  changes  to  a  dark  liquid  so  viscous 
that  the  test  tube  may  be  inverted  without  spilling  the  sulphur.  Heat  to 
boiling  (448.5°);  observe  that  the  fluid  becomes  less  viscous.  Collect  a 
portion  of  the  sulphur  vapor  in  a  cold  flask  or  bottle. 

Pour  the  boiling  sulphur  from  the  test  tube  into  a  beaker  of  water. 
Examine  the  cooled  mass.  Does  it  resemble  roll  sulphur  ?  Is  it  amorphous 
or  crystalline?  It  is  known  as  plastic  sulphur.  Place  a  part  of  it  on  a 
watch  glass  and  set  aside  to  be  examined  after  a  couple  of  days.  Re- 
sults ?  Reserve  the  other  portion  for  the  following  experiment. 

Note. — Plastic  (elastic)  sulphur  and,  in  fact,  all  amorphous  bodies 
are  spoken  of  as  "  supercooled  liquids."  See  Lecture  Notes. 

Experiment  II. — Monoclinic  Sulphur. 

Melt  25-50  grams  of  sulphur  in  a  porcelain  evaporating  dish  or  a 
Hessian  crucible;  cool  until  a  crust  has  barely  formed  over  the  surface; 
perforate  the  crust  with  a  glass  rod,  and  pour  off  the  liquid  portion. 
The  sides  and  bottom  of  the  dish  will  be  covered  with  transparent,  needle- 
like  crystals  of  sulphur.  Set  the  monoclinic  crystals  aside  for  several 
days,  then  observe  that  they  have  become  opaque  like  roll  sulphur. 

Experiment  III. — Rhombic  Sulphur.  Solubility  of  the  Allotropic  Forms 
of  Sulphur. 

(a)  Place  a  small  piece  of  roll  sulphur  in  the  bottom  of  a  test  tube;  add 
5  or  6  cm.3  of  CS2;  after  shaking  for  some  time,  pour  a  cm.3  of  the  liquid 
upon  a  glass  ;plate;    the    CS2  will  evaporate  spontaneously  depositing 
crystals  of  rhombic  (octahedral)  sulphur.     Examine  the  crystals  with  a 
hand- magnifying  glass  or  by  the  aid  of  a  microscope.     Are  the  crystals 
soluble  in  H2O  ?     In  CS2  ? 

(b)  Repeat  (a)  using  plastic  amorphous  sulphur,  monoclinic   (pris- 
matic) sulphur,   and  flowers  of  sulphur.     Test  each  form  separately. 
Record  results.     Sulphur   crystallizes   in   how   many  forms?     Such    a 
substance  is  given  what  name? 


178  EXPERIMENTAL  CHEMISTRY. 

Native  sulphur  is  a  yellow  crystalline  substance  composed  largely  of 
partially  formed  rhombic  crystals.  It  has  a  sp.  gr.  of  2.06;  melts  at 
114.5°,  and  boils  at  448.5°.  Below  96°,  prismatic  or  monoclinic  sulphur 
is  unstable  and  passes  into  the  rhombic  variety;  above  96°  rhombic 
sulphur  is  unstable  and  passes  into  the  monoclinic.  The  range  of  stable 
rhombic  sulphur  extends  from  low  temperature  to  96°  C.;  that  of  mono- 
clinic  from  96°  to  119.5°,  its  melting-point.  Amorphous  sulphur, 
likewise  monoclinic,  if  allowed  to  stand  at  ordinary  temperatures  becomes 
hard  and  brittle;  it  passes  into  the  rhombic. 

That  which  is  of  special  interest  is  the  fact  that  monoclinic  sulphur 
with  a  sp.  gr.  of  1.96  and  a  melting-point  of  119.5°,  possesses  properties 
quite  different  from  those  possessed  by  the  rhombic  form.  (The  opaque- 
ness which  was  attendant  upon  the  transformation  of  the  monoclinic 
into  rhombic  crystals,  is  due  to  the  fact  that  the  rhombic  crystals  with 
their  greater  sp.  gr.  (2.07)  occupy  less  actual  space  than  the  monoclinic 
or  prismatic.)  The  difference  between  these  two  varieties  of  sulphur 
suggests  that  which  was  observed  between  the  two  forms  of  oxygen, 
namely,  ordinary  oxygen  and  ozone.  (Monoclinic  sulphur  is  frequently 
referred  to  as  the  analogue  of  ozone.)  It  is  quite  doubtful,  however, 
whether  an  analogous  explanation  will  satisfactorily  meet  the  require- 
ments of  both  cases.  It  seems  quite  certain  in  the  case  of  sulphur  that 
the  different  energy  contents  of  the  two  forms  are  in  some  way  intimately 
related  to  the  difference  in  properties.  The  change  from  the  monoclinic 
to  the  rhombic  form  is  accompanied  by  an  evolution  of  a  considerable 
amount  of  heat.  (Compare  ozone.)  It  is  evident  that  the  molecule  of 
monoclinic  sulphur  contains  more  internal  energy  than  the  molecule  of 
rhombic  sulphur.  Favre  and  Silbermann  (Ann.  Chim.  Phys.)  treated 
this,  and  found  the  following: 

S(monoclinic)  +  O2 — »  SO2  +  73,300  cal. 
S  (rhombic)       +  O2  — •*  SO2  +  71,000  cal. 

The  difference,  2300  cal.,  is  the  thermal  equivalent  of  the  difference 
between  the  energy  contents  of  the  two  forms  of  sulphur. 

Experiment  IV. — Union  of  Sulphur  with  Metals  to  Form  Sulphides. 

Place  a  small  globule  of  mercury  in  a  clean  mortar  containing  a  little 
sulphur;  rub  the  two  substances  together  by  use  of  pestle.  What  is  the 
color  of  the  resulting  powder?  Compare  it  with  the  mercury  sulphide 
found  on  the  end-shelf. 

With  the  exception  of  platinum  and  gold,  finely  divided  metals  when 
rubbed  with  sulphur  combine  with  the  latter  to  form  sulphides. 

Experiment  V. — Formation  of  Ferrous  Sulphide.     Hydrogen  Sulphide. 

To  5  grams  of  powdered  sulphur  in  a  mortar,  add  7  grams  of  fine  iron 
filings;  mix  thoroughly  and  introduce  mixture  into  a  small  test  tube;  heat 
the  tube  until  the  contents  glow;  cool,  and  break  the  tube;  examine  the 


OXYGEN   FAMILY.  179 

fused  mass.  Is  it  a  mixture  or  a  compound  ?  To  several  small  pieces  in  a 
test  tube,  add  dilute  HC1  or  H2SO4.  Note  the  odor  of  the  evolved  gas, 
hydrogen  sulphide  or,  as  it  is  sometimes  styled,  sulphuretted  hydrogen. 
Indicate  the  chemical  reactions  by  means  of  equations. 

H2  -f  S  (rhombic)  — »  H2S  +  2730  cals. 

Experiment  VI. — Preparation  and  Properties  of  Hydrogen  Sulphide. 
Hydrosulphuric  Acid. 

Place  several  pieces  of  ferrous  sulphide  (iron  sulphide)  in  the  flask 
of  the  apparatus  used  for  generating  hydrogen;  insert  the  stopper; 
through  the  thistle  tube  introduce  2  cm.3  of  strong  HC1  or  H2SO4.  Ob- 
serve the  slow  evolution  of  the  hydrogen  sulphide,  H2S;  add  H2O  slowly 
until  the  evolution  of  gas  becomes  rapid.  Explain  action  of  H2O. 
Collect  some  of  the  gas  over  water.  Note  its  color,  odor  and  its  density 
as  compared  with  air.  Inflame  a  bottle  of  the  gas.  Explain  the  deposit 
on  the  inside  of  the  test  tube. 

Determine  the  solubility  of  the  gas  in  water.  Test  the  action  of  an 
aqueous  solution  of  it  upon  litmus  paper,  and  a  solution  of  sodium  car- 
bonate. A  water  solution  of  H2S  is  frequently  termed  hydro  sulphuric 
acid  or  sulphrydic  acid. 

Raise  the  delivery  tube  from  the  water;  attach  a  jet  and  ignite  the 
issuing  gas  if  all  the  air  has  been  expelled  from  the  flask.  Hold  a  cold 
piece  of  porcelain — for  example,  an  evaporating  dish  or  a  mortar — in  the 
darker  portion  of  the  flame.  What  substance,  evidently  in  the  free  state 
in  the  inner  portion  of  the  flame,  is  deposited  upon  the  porcelain  ?  What 
other  elementary  substance  probably  exists  in  a  free  state  in  the  interior 
of  the  flame?  State  the  probable  reason  for  the  existence  in  the  free 
state  of  these  two  elementary  substances. 

Note. — Use  the  remaining  gas  to  saturate  50  cm.3  of  H2O  in  a  flask. 
This  aqueous  solution  may  be  used  for  the  following  experiments  if  the 
laboratory  is  not  equipped  with  a  generator  which  furnishes  a  constant 
supply  of  H2S. 

(H2S.  Aq)  =  4500  cal. 
Experiment  VII. — Properties  of  Hydrogen  Sulphide. — Continued. 

(a)  Instability  of  hydrogen  sulphide.     In  a  clean  test  tube  expose  an 
aqueous  solution  of  the  gas  to  the  action  of  the  air  for  several  days. 
Results  ?     Equation  ? 

(b)  Reducing  action  of  hydrogen  sulphide.     Pass  a  stream  of  H2S 
through  5  cm.3  of  strong  H2SO4-  until  there  is  evidence  that  SO2  is  being 
evolved  and  free  sulphur  is  deposited.     Equation?     Could  H2SO4  be 
used  to  dry  H2S  ? 

To  5  cm.3  of  a  K2Cr2O7  (potassium  dichromate)  solution  add  2  cm.s 
of  HC1.  What  is  the  color  of  the  solution?  Pass  H2S  through  the 
solution  until  its  color  is  green,  due  to  the  presence  of  CrCl3.  Is  free 


l8o  EXPERIMENTAL  CHEMISTRY. 

sulphur  deposited?  Write  the  structural  formula  for  K2Cr2O7.  What 
is  the  valence  of  the  chromium  in  this  salt?  Structural  formula  for 
CrCl3  ?  Valence  of  chromium  in  this  salt  ?  Write  equations  showing 
nature  of  the  reactions. 

The  above  may  be  repeated  using  KMnO4  instead  of  K2Cr2O7. 

To  3  cm.3  of  ferric  chloride,  FeCL3  add  a  few  drops  of  potassium 
ferricyanide,  K3Fe(CN)6.  Note  the  color.  Repeat,  using  ferrous 
sulphate,  FeSO4,  instead  of  FeCl3.  Note  the  color.  Pass  H2S  through 
separate  solutions  of  FeCl3,  and  FeSO4,  for  two  or  three  minutes;  filter 
and  test  each  filtrate  with  K3Fe(CN)6.  Explain  results.  Equations? 

(c)  Formation  of  sulphides.     Allow  hydrogen  sulphide  to  come  into 
contact  with  a  moist  silver  coin.     Results  ?     Equation  ? 

(d)  Determine  the  electrical  conductivity  of  an  aqueous  solution  of 
H2S.     What  is  your  conclusion  as  to  its  degree  of  dissociation  ?     Is  it  a 
strong  or  weak  acid  ?     Why  ? 

Experiment  VIII. — Hydrogen  Sulphide  and  Sulphides. 

Introduce  3  cm.3  of  ZnSO4  solution  into  a  test  tube  add  5  cm.3  of 
water.  Test  the  solution  with  litmus  paper.  Is  the  solution  neutral? 
Saturate  solution  with  H2S.  Does  a  precipitate  of  ZnS  form?  Is  the 
precipitate  soluble  in  HC1  ?  Test  by  adding  2  cm.3  of  said  acid.  Saturate 
solution  a  second  time  with  the  H2S.  Is  a  precipitate  formed?  Write 
the  equations  for  the  above  reactions  showing  the  interaction  of  the  ions. 

Repeat  the  foregoing,  using  solutions  of  CuSO4,  Cd(NO3)2,  NaCl, 
MnSO4,  BaCl2,  H3AsO3,  in  place  of  ZnSO4.  Classify  the  sulphides 
according  to  their  solubility  in  water,  dilute  acids  and  strong  acids. 

Sulphides  of  the  alkalies  and  the  alkaline  earths  are  soluble  in  water; 
all  other  metallic  sulphides  are  insoluble.  The  behavior  of  the  insoluble 
sulphides  towards  acids  may  be  summarized  as  follows: 

(a)  Sulphides  decomposed  by  dilute  acids  like  HC1  or  H2SO4.     Ex- 
amples:    FeS,  ZnS,  MnS. 

(b)  Sulphides  not  decomposed  by  dilute  acids,  but  soluble  in  warm 
strong  HC1.     Examples:     PbS,  Sb2S3,  SnS,  NiS,  CoS. 

(c)  Sulphides  not  soluble  in  strong  HC1,  but  are  decomposed  by  aqua 
regia  or  by  a  mixture  of  HC1  and  KC1O3.     Examples:     HgS,  As2S3. 

Experiment  IX. — A  Test  for  Sulphides. 

(a)  When  soluble  sulphides  are  decomposed  by  acids,  HjS  is  evolved. 
The  presence  of  the  gas  may  be  detected  (i)  by  its  odor;  (2)  by  its  action 
upon  lead  acetate.     This  latter  test  is  usually  made  as  follows:     A 
piece  of  filter  paper  moistened  with  lead  acetate  is  held  over  the  mouth 
of  the  test  tube  in  which  the  reaction  occurs.     The  H2S  causes  a  black 
stain  due  to  the  formation  of  lead  sulphide. 

(b)  Prepare  a  dilute  solution  of  Na,jS  or  K2S.     Note  that  the  odor  of 
H2S  is  not  present.     Following  (a),  test  for  sulphides.     Equations  ? 


OXYGEN   FAMILY.  l8l 

Experiment  X. — Optional.     Polysulphides.     Hydrogen  Poly  sulphide. 

(a)  To  20  cm.3  of  a  solution  of  a  soluble  sulphide,  e.g.,  Na^S,  add  a 
pinch  of  pulverized  sulphur;  shake  vigorously  until  the  color  of  the 
solution   becomes   reddish-yellow.     After  evaporation   of   the   solution, 
the  substances  remaining  have  a  composition  varying  from  Na^  to 
Na2S5.     Dissolve  the  residue  in  water.     Divide  the  solution  into  two 
parts.     To  one  portion  add  strong  HC1  and  note  the  liberation  of  H2S 
and  free  sulphur  (rhombic).     Add  the  other  portion  to  3  cm.3  of  strong 
HC1.     Observe  that  no  H2S  is  evolved,  but  that  a  heavy  yellow  oil, 
hydrogen  pentasulphide  (H2S5)  forms  a  distinct  phase  in  the  bottom  of 
the  tube.     H2S5  is  unstable  in  the  presence  of  moisture  and  decomposes 
into  H2S  and  S. 

(b)  Add  a  solution  of   CaS2,  (CaSx)  to  HC1;   the  yellow  oily  liquid, 
hydrogen  polysulphide,  separates. 

Oxygen  Derivatives  of  Sulphur. 

Experiment  XI. — Formation  of  Sulphur  Dioxide. 

(a)  Oxidation  of  sulphur  by  burning  in  the  air.     Set  fire  to  a  few 
small  pieces  of  sulphur  upon  the  inverted  cover  of  a  porcelain  crucible. 
Observe  the  color  of  the  flame  and  note  the  odor  of  the  sulphur  dioxide, 
SO2,  fumes. 

(b)  Roasting  a  pyrite,  FeS2.     Heat  i  gram  of  granular  pyrite  in  a 
hard  glass  test  tube.      Observe   the   deposit  of  sulphur  on   the   cooler 
portions  of  the  tube, and  notice  the  odor  of  evolved  SO2.     What  is  the 
effect  of  heating  those  portions  of  the  tube  on  which  sulphur  is  deposited  ? 

(c)  Reduction   of   H2SO4   by  sulphur.     Heat   a   small   piece   of   roll 
sulphur  with  strong  H2SO4  in  a  test  tube.     Wave  the  escaping  fumes 
toward  the  nose.     Odor?     Equation?    Pieces  of  charcoal  may  be  substi- 
tuted for  the  sulphur,  and  the  above  procedure  followed. 

(d)  Reduction  of  H2SO4  by  copper.     See  Exp.  XII. 

S  +  O2  — »  SO2  +  71,000  cal. 

Would  you  infer  from  equation  that  SO2  is  a  stable  or  an  unstable 
compound  ? 

Experiment  XII. — (Hood)  Laboratory  Preparation  of  Sulphur  Dioxide. 

Fit  a  300  cm.3  flask  with  a  rubber  stopper  provided  with  thistle  tube 
and  a  delivery  tube,  and  add  two  or  three  small  bunches  of  copper 
shavings  or  10  grams  of  granulated  copper;  support  the  flask  in  ring 
stand  upon  a  wire  gauze;  connect  flask  (Fig.  29)  with  wash  bottle,  one- 
third  filled  with  strong  H2SO4;  the  wash  bottle  should  be  provided  with 
a  doubly  bent  delivery  tube,  reaching  almost  to  the  desk-top.  Add  through 
thistle  tube  100  cm.3  to  150  cm.3  of  strong  H2SO4;  heat  bottom  of  flask 
gently  until  there  is  a  brisk  evolution  of  gas.  Identify  the  gas.  Fill  three 
bottles  with  the  gas  by  downward  displacement  of  air;  stopper  the  bottles. 
Saturate  25  cm.3  of  H2O  with  the  gas;  the  delivery  tube  should  be  placed 


i82 


EXPERIMENTAL  CHEMISTRY. 


i  cm.  above  the  surface  of  the  water.  What  evidence  is  there  that  the 
gas  is  dissolving?  Proceed  with  the  study  of  the  properties  of  SO2. 
Equation  indicating  action  of  strong  H2SO4  on  Cu? 

Experiment  XIII. — Properties  of  Sulphur  Dioxide. 

(a)  Bleaching  properties  of  sulphur  dioxide.  Place  a  moist  strip 
of  colored  calico  and  the  moistened  petals  of  a  red  flower  in  one  bottle 
of  the  gas.  Set  them  aside  until  the  close  of  the  laboratory  period. 
Effect  of  gas? 


FIG.  29. 

(b)  Using  another  bottle  of  the'gas,  ascertain  if  SO2  will  burn  or  support 
combustion.     Conclusions  ? 

(c)  Devise  a  method  for  determining  the  density  of  SO2  as  compared 
with  air.     Record  procedure  and  results. 

(d)  What  is  the  effect  of  dry  SO2  upon  dry  blue  litmus  paper? 

Experiment  XIV. — Optional.     Liquid  Sulphur  Dioxide. 

Note. — The  critical  temperature  of  SO2  is  156°  C.     The  gas,  SO2, 
when  cooled  to  the  temperature  of  a  freezing  mixture  of  salt  and  ice, 
condenses  to  a  liquid  which  boils  at  about  — 8°  C. 

Pass  a  slow  current  of  dry  SO2  through  a  long  spinal  tube,  immersed 
in  a  freezing  mixture,  into  a  U-tube,  the  arms  of  which  are  provided 
with  glass  stopcocks.  This  latter  tube,  "sulphur  dioxide  condenser," 
must  also  be  placed  in  a  freezing  mixture.  The  liquid  SO2  may  be 
kept  indefinitely  in  the  closed  tube.  When  gaseous  SO2  is  required, 
one  of  the  stop  cocks  is  opened. 


OXYGEN    FAMILY.  183 

Experiment  XV. — Aqueous  Solution  of  Sulphur  Dioxide.  Sulphurous 
Acid. 

Test  the  action  of  the  aqueous  solution  of  SO2  prepared  in  Exp.  XII. 
upon  blue  litmus  paper.  To  i  cm.3  of  the  solution  add  10  cm.3  of  H2O. 
Taste  a  drop  of  the  solution.  Does  the  solution  manifest  the  properties 
of  an  acid,  base  or  salt?  Explain  its  formation.  Equation?  Heat 
5  cm. 3  of  the  acid  solution.  Note  the  odor  of  the  fumes.  SO2  is  the 
anhydride  of  what  acid?  What  is  the  significance  of  the  ous  ending? 
What  is  the  termination  of  its  corresponding  salts  ?  Write  the  structural 
formula  of  sulphurous  acid.  How  many  series  of  salts  can  H2SO3 
form  ?  What  name  is  applied  to  such  acids  ?  Would  you  say  that 
H2SO3  is  a  very  stable  compound?  Why? 

SO2  +  H2O  —  SO2,  Aq  +  7,700  cal. 

Experiment  XVI. — Preparation  of  a  Soluble  Sulphite.  Action  of  Strong 
Acids  Upon  Soluble  Sulphites.  Test  for  Sulphites. 

Place  10  cm.3  of  H2SO3  in  a  test  tube;  add  a  solution  of  NaOH  until 
the  odor  of  SO2  has  disappeared.  Divide  the  sodium  sulphite  solution 
into  two  parts.  To  one  portion  add  a  little  HC1;  to  the  other  portion 
add  H2SO4.  Is  SO2  evolved  in  both  cases?  Equations? 

Experiment  XVII. — A  Test  for  Sulphites.     Reversible  Reaction. 

To  5  cm.3  of  a  BaCl2  solution  add  drop  by  drop  several  cm.3  of  H2SO3. 
Observe  that  the  white  precipitate  of  BaSO3  is  not  heavy.  Write  the 
equation  for  the  reaction.  Is  the  precipitate  soluble  in  HC1?  What 
acid  was  formed  when  BaSO3  was  precipitated?  Will  this  acid  have  a 
tendency  to  reverse  the  action  of  the  H2SO3?  Is  it  probable  that  the 
precipitation  of  BaSO3  would  be  more  complete  if  the  acid  were  removed 
as  fast  as  it  formed  ? 

Ascertain  the  effect  of  adding  a  solution  of  sodium  or  potassium 
acetate  to  a  solution  of  BaCl3  to  which  has  been  added  H2SO3.  Is  the 
precipitation  more  complete  ?  Is  the  precipitate  soluble  in  HC1?  What 
gas  is  evolved?  Explain  the  further  precipitation  of  BaSO3  on  the 
addition  of  sodium  acetate,  NaC2H3O2,  by  use  of  " ionic"  equations. 

All  sulphites  except  those  of  the  alkali  metals  are  insoluble  in  water. 
Dilute  acids  decompose  all  sulphites  with  an  evolution  of  SO2.  The 
larger  number  of  the  sulphites  are  converted  into  sulphides  and  sulphates 
when  heated.  Those  of  the  alkaline  earths  yield  a  metallic  oxide  and 
S02. 

4Na2SO3->Na2S  +  3Na,SO4, 
BaSO3->BaO  +  SO2. 

Experiment  XVIII. — Reducing  Action  of  Sulphurous  Acid. 

(a)  What  is  the  effect  of  adding  H2SO3  to  a  very  dilute  solution  of 
KMnO4  acidulated  with  H2SO4?  Equations? 


184  EXPERIMENTAL  CHEMISTRY. 

(b)  Substituting  K2Cr2O7  for  KMnO4,  repeat  (a).     Equations? 

(c)  Recall  action  of  H2SO3  on  HIO3.     Was  reduction  accompanied 
by  oxidation  ?     Explain. 

Sulphurous  acid  in  the  presence  of  more  powerful  reducing  agents 
may  act  as  an  oxidizing  substance  as  follows: 

Experiment  XIX.  —  Oxidizing  Action  of  Sulphurous  Acid. 

(a)  Add  2  cm.3  of  HC1  to  5  cm.3  of  SnCl2;  add  2  cm.3  of  H2SO3;  warm 
gently;  observe  the  formation  of  a  yellow  precipitate  of  SnS2.  As  the 
reactions  usually  proceed  rather  slowly,  allow  the  test  tube  and  contents 
to  remain  undisturbed  for  10  minutes  to  15  minutes.  The  H2SO3  is 
reduced  to  H2S,  and  the  SnCl2  is  oxidized  to  SnCl4. 


(1)  3SnCl2  +  6HC1  +  H2SO3  —  3SnCl4  +  3H2O  +  H^Si 

(2)  SnCl4  +  2H2S      —  SnS2  +  4HC1. 

(b)  Nascent  hydrogen  reduces  sulphurous  acid.  Place  a  piece  of 
granulated  zinc  in  the  bottom  of  a  test  tube;  add  a  little  dilute  HC1;  add  a 
few  drops  of  H2SO3.  Test  for  the  presence  of  H2S  by  holding  a  piece  of 
filter  paper  moistened  with  a  drop  of  lead  acetate  over  the  mouth  of 
the  tube  (?).  Equations? 

Experiment  XX.  —  Oxidation  of  Sulphurous  Acid.     Sulphuric  Acid. 

(a)  To  3  cm.3  of  H2SO3  add  a  little  H2O2;  shake,  then  add  a  small 
quantity  of  a  BaCl2  solution.     Is  the  precipitate  which  forms  soluble  in 
HC1?     Recall  the  solubility  of  BaSO3  in  HC1.     The  white  precipitate 
insoluble  in  HC1  is  BaSO4.     What  acid  was  evidently  present  in  the 
solution    before    the    BaCl2    was    introduced?     Explain    its   formation. 
Equations?     Using  H2SO4  from  the  reagent  shelf,   test  its  action  on 
BaCl2.     Is  the  precipitate  soluble  in  HC1?     How  can  you  distinguish 
between  a  sulphite  and  a  sulphate  ? 

(b)  Repeat  (a)  using  "bromine  water."     Results?     Equations? 

(c)  Slow  oxidation  of  H2SO3  by  oxygen  of  the  air.     Expose  10  cm.3 
of  H2SO3  to  the  action  of  the  air  for  several  hours.     Test  it  with  BaCl2. 
Was  BaSO3  or  BaSO4  formed  ?     Give  reasons  for  answer.     Equations  ? 

(b)  Rapid  oxidation  of  H2SO3  by  oxygen  of  the  air  by  use  of  a  catalyser, 
N2O3.  Prepare  the  gaseous  catalyser,  N2O3,  as  follows:  Put  10  or 
15  grams  of  granulated  copper  or  a  bunch  of  copper  shavings  in  a 
generating  flask;  add  sufficient  water  through  the  thistle-  tube  to  cover 
the  lower  end  of  the  latter;  then  introduce  20  cm.3  of  concentrated  HNO3. 
Observe  the  formation  of  a  reddish-brown  gas  which  is  a  mixture  of  N2O3 
and  N2O4.  The  colorless  gas,  NO,  is  the  initial  product  formed  by  the 
action  of  HNO3  on  copper,  but  in  the  presence  of  air,  it  quickly  oxidizes 
to  the  above-mentioned  gases.  Fill  a  small  flask  with  NO  by  displace- 
ment of  water;  allow  air  to  enter  flask  that  N2O3  may  be  formed,  then 
pour  2  cm.3  of  H2SO3  into  the  flask;  fit  a  stopper  in  place  and  shake. 


OXYGEN    FAMILY. 


Observe  that  the  color  disappears.  Remove  stopper  until  gas  is  again 
colored;  insert  stopper  and  shake  (?).  Repeat  this  operation  several 
times,  then  by  use  of  BaCl2  and  HC1  prove  the  presence  of  H2SO4. 


Experiment  XXI. — (L.  T.) 
Chamber  Process. 


The  Manufacture    of    Sulphuric  Acid. — 


The  purpose  of  the  following  experiment  is  to  acquaint  the  student 
with  the  "chemistry"  of  the  sulphuric  acid  industry.  References  and 
lectures  should  supplement  the  experiment.  A  flask  or  glass  globe  of  two 
or  three  liters'  capacity  is  fitted  with  a  cork  through  which  passes  five  de- 
livery tubes  reaching  to  the  center  of  the  flask  (Fig.  30).  The  middle  tube 
e  serves  as  an  escape  for  the  excess  of  gases  introduced  through  the  other 


FIG.  30. 

tubes.  Through  a  tube  d  air  can  be  forced  into  the  flask  (chamber) 
by  a  bellows  or  a  bicycle  pump.  The  other  three  tubes  are  connected 
by  glass  tubing  with  separate  generating  flasks  a,  b  and  c.  a  is  the  flask 
in  which  SO2  is  prepared  by  heating  copper  shavings  and  concentrated 
H2SO4.  (The  SO2  may  be  prepared  by  heating  granular  pyrite  in  a 
boat  placed  in  hard  glass  tubing.  A  current  of  air  should  be  drawn 
through  the  tube  over  the  pyrite  by  use  of  an  aspirating  bottle.)  Nitric 
oxide,  N2O2,  which  in  the  presence  of  oxygen  oxidizes  to  N2O3,  the  cataly- 
ser,  is  generated  in  b  by  the  action  of  dilute  HNO3  on  copper  shavings. 
(Nitric  acid  vapor  may  be  used  instead.)  Steam,  H2O,  is  supplied  by 
boiling  water  in  flask  c.  If  the  products  (gases)  from  a  and  b  together 
with  air  are  introduced  without  a  sufficient  supply  of  steam,  a  frost-like 


1 86  EXPERIMENTAL  CHEMISTRY. 

deposition  of  nitrosyl-sulphuric  acid  crystals  is  observed  upon  the  walls 
of  the  flask. 

H  O  — H. 

\  / 

O  +  2SO2  +  N2O3  +  O2— >  2SO2 

/  \ 

H  O  — NO. 

Now  increase  the  supply  of  steam  and  note  the  disappearance  of  the 
crystals  and  the  formation  of  a  heavy  oily  fluid: 

O  — H  OH 

/  / 

2S02  +  H20  *±  2S02  +  N203. 

\  \ 

O  — NO  OH 

When  10  cm.3  to  15  cm.3  of  liquid  have  collected  in  the  flask,  the 
experiment  may  be  interrupted.  Place  the  liquid  in  a  test  tube  and 
reserve  it  for  following  experiments. 

Experiment  XXII. — Properties  of  Sulphuric  Acid  (Hydrogen  Sulphate). 

(a)  Dehydrating  action  of  sulphuric  acid.     Action  of  sulphuric  acid 
with  organic  matter.     Remove  the  head  of  a  match;  dip  the  \vood  into 
hot    concentrated    H2SO4.     Results?     Write    your    name    with    dilute 
H2SO4  upon  a  piece  of  paper.     Dry  the  paper.     Results? 

To  a  little  sugar  in  the  bottom  of  a  test  tube  add  strong  H2SO4;  heat 
gently.  Results?  Each  of  the  above  substances  is  composed  largely 
of  C,  H,  and  O.  What  is  the  dark-colored  substance  which  remains  in 
each  case  ? 

Repeat  each  of  the  above  experiments  with  the  oily  fluid  prepared  in 
Exp.  XXI.  Make  a  record  of  results.  Are  its  properties  identical 
with  the  shelf-reagent,  H2SO4  ?  Is  sulphuric  acid  a  efficient  drying  agent  ? 

(b)  To  i  cm.3  of  the  fluid  prepared  in  Exp.  XXI,  add  5  cm.3  of  water. 
Test  with  blue  litmus  paper.     Repeat  above,  using  H2SO4.     Results? 

(c)  Try  the  action  of  the  "acid"  on  a  piece  of  zinc  in  a  test  tube. 
What  gas  is  evolved?     Repeat  with  H2SO4.     Compare  results. 

(d)  Add  BaCl2  to  a  portion  of  the  acid  prepared  in  above  experiment. 
Is  the  precipitate  soluble  in  HC1? 

Does  H2SO4  yield  similar  results?  Drawing  your  inference  from 
the  results  of  the  above  experiments,  identify  the  liquid  prepared  in 
Exp.  XXI.  Write  the  structural  formula  for  H2SO4. 

Experiment  XXIII. — Hydration  of  Sulphuric  Acid. 

(a)  Into  5  cm.3  of  water  in  a  test  tube  add  10  cm.3  of  strong  H2SO4. 
Is  there  a  change  in  the  temperature  of  the  liquids  ?  What  is  the  probable 
explanation  of  the  thermal  phenomenon? 

H2SO4+  2H2O  — H2SO4 ,  Aq  +  17,850  cal. 

Is  the  solution  more  or  less  stable  than  the  pure  acid  ?     Why  ? 


OXYGEN    FAMILY.  1 87 

(b)  Optional.  The  following  interesting  experiment  is  suggested  by 
Freer:  Add  9.8  grams  of  concentrated  H2SO4  to  1.8  grams  of  water;  place 
the  liquid  in  a  small  flask  and  surround  the  latter  with  a  freezing  mixture 
of  ice  and  salt;  crystals  of  H2SO4.H2O  (H4SO5)  will  form;  warm  the 
crystals  until  they  are  melted  (80°  C.),  then  add  1.8  grams  of  water;  again 
place  in  the  freezing  mixture  and  crystals  of  H2SO4.2H2O  (H6SO6)  will 
separate.  Further  addition  of  water  is  not  accompanied  by  increase 
of  temperature. 

H2SO4,  Normal  Sulphuric  Acid. 

/OH    Ox      /  O  —  H 
H2S04,  S02(OH)2,  S02  (  >S( 

XOH,  O/    X0  —  H. 

H2SO4.H2O,  Tetra  hydroxyl  Sulphuric  Acid. 

/OH  /OH 

"/OH  //OH 

H4S05,     SO/        ,  O  =  S/         , 

XOH  XOH 

XOH  XOH 

H2SO4.2H2O,  Hexahydroxyl  Sulphuric  Acid. 

/OH 
/OH 

/OH 
H6S06,         S(OH)6,          S< 

XOH 
XOH 
XOH 

Experiment  XXIV. — Dissociation  of  Sulphuric  Acid. 

Pour  4  cm.3  of  concentrated  H2SO4  into  a  clean  dry  test  tube;  clamp 
the  tube  to  the  ring  stand;  suspend  a  2oo°-3oo°  C.  thermometer  so  that 
the  bulb  will  be  immersed  in  the  acid;  heat  the  test  tube  gently  with  a 
small  flame.  Take  the  reading  on  the  thermometer  at  the  first  appear- 
ance of  the  heavy,  dense  white  fumes  of  SO3.  Caution. — Hot  H2SO4 
produces  severe  burns.  Note  the  thermometer  reading  frequently, 
that  the  thread  of  mercury  may  not  be  allowed  to  approach  too  near 
the  top  of  the  tube.  Write  the  equation  indicating  the  dissociation  of 
H2SO4.  What  is  the  difference  between  dissociation  and  decomposition  ? 

Experiment  XXV. — Reduction  of  Sulphuric  Acid. 

Recall  Exp.  XI.  (r),  (d).     Equations? 

Experiment  XXVI.— Soluble  and  Insoluble  Sulphates. 

Test  the  action  of  a  solution  of  Na2SO4  or  dilute  H2SO4  upon  separate 
solutions  of  each  of  the  following  substances:  AgNO3,  NaCl,  CaCl2, 


1 88  EXPERIMENTAL  CHEMISTRY. 

Pb(N03)2,  Cu(N03)2,  KN03,  BaCl2,  Sr(NO3);  NH4C1,  FeCl3.  Write 
the  ionic  equations  for  each  reaction.  Underscore  the  formula  repre- 
senting the  precipitate  in  each  case. 

Experiment  XXVII. — Optional.     Decomposition  of  Isoluble  Sulphates. 

Insoluble  sulphates,  e.g.,  BaSO4,  may  be  decomposed  by  fusion  with 
Na^CC^;  Na^SC^,  which  is  soluble  in  water,  being  formed.  The  fused 
mass  is  extracted  with  water  and  the  solution  made  slightly  acid  with 
HC1,  after  which  BaCl2  is  added.  The  substances  are  usually  fused  in  a 
crucible. 

Note. — The  above  method  is  quite  generally  used  for  converting 
insoluble  salts  into  corresponding  salts  of  the  alkali  metals.  Practically 
all  of  the  salts  of  sodium  and  potassium  are  soluble  in  water,  hence  the 
use  of  Na-jCOg  or  K2CO3  as  a  flux. 

Most  sulphates  are  soluble  in  water.  Silver,  calcium,  strontium, 
lead  and  barium  are  but  slightly  soluble  in  water.  Their  solubility 
decreases  in  the  order  named. 

Experiment  XXVIII. — Optional.    Reduction  of  Sulphates  and  Sulphites. 

(a)  Mix  a  small  quantity  of  a  sulphate  or  a  sulphite  with  two  or  three 
times  its  weight  of  pure  anhydrous  Na^COg;  place  a  portion  of  the 
mixture  in  a  depression  in  a  piece  of  charcoal;  using  a  blow-pipe,  direct 
the  reducing  flame  upon  the  mixture.  Place  the  fluid  mass  upon  a  silver 
coin  or  a  piece  of  filter  paper  moistened  with  lead  acetate;  allow  a  drop 
of  dilute  HC1  to  fall  upon  the  mass.  Explain  the  formation  of  a  black 
stain.  Equations  ? 

Experiment  XXIX. — Optional.  Sulphur  Trioxide.  The  Manufacture 
of  Sulphuric  Acid  by  the  "Contact  Process." 

(a)  Put   15   cm.3  of  fuming  sulphuric  acid   (disulphuric  or  pyrosul- 
phuric  acid,  H2S2O7)  in  a  small  dry  retort  provided  with  a  glass  stopper 
and  connected  with  a  dry  receiver  immersed  in  water.     Keep  the  re- 
ceiver cool.     Heat  the  flask  gently,  and  observe  evolution  of  the  heavy 
SO3  fumes  which  condense  to  a  solid  in  the  receiver.     Remove  some  of 
the  solid  by  means  of  a  glass  rod  and  put  it  in  a  little  water.     Is  heat 
evolved  ?    Place  other  portions  on  pieces  of  paper  and  wood.     Results  ? 

(b)  Sulphur  dioxide  and  oxygen  do  not  combine  rapidly  under  or- 
dinary conditions  to  form  sulphur  trioxide.     At  higher  temperatures 
the  reaction  proceeds  relatively  slowly.    Above  400  °  C. ,  S  O3  is  decomposed 
into  SO2  and  O.     It  has  been  found  that  if  the  two  gases  are  passed 
simultaneously  over  heated  finely-divided  platinum  they  combine  to  form 
SO3  which  combines  readily  with  water  to  form  H2SO4.     The  above  is  a 
general  statement  of  the  principles  involved  in  the  manufacture  of  H2SO4 
by  the  "contact  process." 

Prepare  platinized  asbestos  as  previously  directed;  place  it  in  a  long 


OXYGEN   FAMILY.  189 

hard  glass  tube.  Devise  a  method  for  passing  oxygen  and  sulphur 
dioxide  through  the  tube.  The  gases  should  be  dry.  Collect  the  gases 
in  a  dry  receiving  flask.  Is  there  any  evidence  of  SO3  being  formed? 
Now  gently  heat  the  tube  beneath  the  asbestos.  Results?  When  the 
products  are  added  to  water,  the  solution  may  be  tested  for  sulphuric 
acid. 

502  +  O      — >  SO3  +  32,100  cal. 

S    '"  +  O3     —  SO3  (liquid)  +  103,200  cal. 

503  +  H2O  —  H2SO4  +  21,300  cal. 
(H2SO4,Aq)  =  17,800  cal. 

(S,  t>8,  Aq)    -       ?      cal. 
(H2,S,04)      =       ?      cal. 

Experiment    XXX. — Preparation  of  a  Soluble  Thiosulphate. 

Dissolve  5  grams  of  sodium  sulphite,  NaSO3  in  20  cm.3  of  H2O,  then 
add  3  grams  of  finely  divided  sulphur,  e.g.,  flowers  of  sulphur.  Boil 
the  mixture  for  10  or  15  minutes.  Replace  the  water  which  has  evapor- 
ated ;  filter.  Pour  the  filtrate  into  a  small  beaker  or  a  crystallizing  dish. 
Sodium  thiosulphate,  Na^Og^ELjO,  commonly  called  hyposulphite 
of  soda  ("hypo"),  will  crystallize  out  of  the  solution.  Acids,  e.g., 
H2SO4,  decompose  thiosulphates  with  the  formation  of  a  sulphate, 
sulphurous  acid  and  the  liberation  of  free  sulphur.  Try  the  action  of 
acids  upon  the  crystals.  Results  ?  Equations  ? 

Experiment  XXXI. — Preparation  of  an  Insoluble  Thiosulphate. 
Decomposition  of  Thiosulphates  by  Heat. 

(a)  To  a  few  cm.3  of  a  solution  of  a  soluble  thiosulphate  add  a  few 
drops  of  a  AgNO3  solution.  Observe  the  color  of  the  precipitate.  Heat 
the  mixture  and  note  the  formation  of  black-silver  sulphide.  Is  SO2 
evolved  ?  Is  free  sulphur  liberated  ?  Equations  ? 

(6)  Lead  acetate  may  be  substituted  for  AgNO3  and  (a)  repeated. 

Experiment  XXXII. — Oxidation  of  Sodium  Thiosulphate  to  a  Tetra- 
thionate  by  the  Action  of  Iodine. 

Prepare  a  solution  of  sodium  thiosulphate;  add  a  little  starch  water. 
Dissolve  a  small  quantity  of  iodine  in  a  potassium  iodide  solution.  To 
5  cm.3  of  the  thiosulphate  solution  add  the  solution  of  iodine  drop  by 
drop  until  the  solution  turns  blue.  Explain.  Equations? 

Note. — The  titration  of  iodine  against  sodium  thiosulphate  with  starch 
as  an  indicator  is  a  process  frequently  used  in  analytical  chemistry. 

Experiment  XXXIII. — (Quant.)  Compare  the  electrical  conductivi- 
ties of  pure  concentrated  H2SO4  and  a  5N  solution  of  same  acid.  Ex- 
plain results.  The  activity  of  an  acid  depends  on  what?  Would  you 


1 90  EXPERIMENTAL  CHEMISTRY. 

infer  that  H2SO4  in  aqueous  solution  is  active  as  an  acid?  Are  the  elec- 
trical conductivities  of  acids  related  in  any  way  to  their  activities  ?  Ex- 
plain. Write  the  ionic  equation  indicating  the  different  particles  present 
in  an  aqueous  solution  of  sulphuric  acid. 

Experiment  XXXIV.— (Quant.)     Specific  Gravity  of  Sulphuric  Acid. 

Devise  a  method  for  determining  the  specific  gravity  of  sulphuric 
acid.  What  is  the  sp.  gr.  of  the  concentrated  H2SO4  which  is  found 
among  the  shelf-reagents  ? 

Experiment  XXXV.— (Quant.)  Contraction  in » Volume  When  Sul- 
phuric Acid  and  Water  are  Mixed. 

By  means  of  a  pipette  place  10-20  cm.3  of  concentrated  H2SO4  in  a 
long  narrow  test  tube  or  a  colorimetric  tube;  carefully  add  one-half  its 
volume  of  water;  mark  the  volume  occupied  by  the  two  liquids.  Now 
mix  the  two  liquids  by  careful  manipulation  of  the  tube.  Note  the 
change  in  temperature  of  the  liquids.  Cool  the  mixture  and  determine 
the  amount  of  contraction.  Does  the  acid  undergo  hydration  ? 

Although  several  factors  may  contribute  to  the  contraction,  it  is  known 
that  "  water  of  hydration"  occupies  less  space  than  ordinary  water. 

OXACIDS    OF    SULPHUR. 

H2SO3  Sulphurous  acid. 

H2SO4  Sulphuric  acid. 

H2S2O3  Thiosulphuric  acid. 

H2S2O4  Hyposulphurous  acid. 

H2S2O5  Pyrosulphurous  acid. 

H2S2O6  Dithionic  acid. 

H2S2O7  Pyrosulphuric  acid. 

H2S2O8  Persulphuric  acid. 

H2S3O6  Trithonic  acid. 

H2S4O6  Tetrathionic  acid. 

H2S5O6  Pentathionic  acid. 

H2S6O6  Hexathionic  acid. 


SELENIUM,    SC. 

At.  Wt.  79.2     Mol.  Wt.  158.4. 
See  Text-book  and  Lecture  Notes. 

TELLURIUM,    T6. 

At.  Wt.  127.6.     Mol.  Wt.  (?). 
See  Text-book  and  Lecture  Notes. 


OXYGEN   FAMILY. 


The  following  table  gives  a  general  view  of  the  similarity  of  the  ele- 
ments of  this  group: 

Oxygen          Sulphur  Selenium  Tellurium 
16.0              32.06  79.2  127.6 
Gas              Solid  Solid  Solid 
Colorless,        Pale-  Grayish-  Silver- 
Yellow  White  White 


Physical  Properties 
Atomic  weight 
State  or  phase 
Color 


Specific  gravity 

I-I3 

2.02 

4-5 

6.3? 

(at—  182°) 

(Rhombic) 

(Crystalline) 

Specific  heat 

s 

^ 

Melting-point 

Below—  223°  111.5° 

217° 

452° 

Boiling-point 

-182° 

445° 

675° 

1400° 

Chemical  Properties 

Heat  of  formation 

H2O 

H2S 

H2Se 

H2Te 

of  H-  derivatives     57 

,000  cal.    2, 

730  cal. 

-5,400  cal. 

—19,400  cal. 

Stability  of  H- 

HATTVQ  fi"\7rpkc 

^s 

<^ 

LLCllVctllVCo 

Heat  of  formation 

02,0(gas) 

S02(gas) 

SeO2(solid) 

TeO2(solid) 

of  O-  derivatives        32,400  cal. 

71,000  cal. 

57,700  cal. 

77,000  cal. 

Oxacids 

H2SO3 

H2S303 

H2Te03 

? 

(solid) 

(solid) 

H2SO4 

H2Se04 

H2TeO4 

(liquid) 

(liquid) 

(liquid) 

Heat  of  formation 

0,CL 

S2C12 

Se2Cl2 

TeCl4 

of  Cl-derivatives        —18,000  cal. 

14,300  cal. 

22,200  cal. 

77,400  cal. 

Metallic  Properties 

_ 

(Physical) 

PROBLEMS. 

1.  How  many  grams  of  sulphur  in  il.  of  H2SO4,  sp.  gr.  1.84?     How 
many  grams  of  oxygen?     Of  hydrogen? 

2.  How  many  grams  of  H2SO4  will  be  required  to  neutralize  10  grams 
ofBa(OH)2? 

3.  How  many  cm.3   of   a    5N  solution  of  H2SO4  will  be  required  to 
neutralize  50  cm.3  of  iN  of  NaOH? 

4.  Calculate  the  heat  of  formation  of  anhydrous  sulphuric  acid  from 
its  elements,  i.e.,  (S,  O4,  H2)  =  ?  Ans.  192,900  cal. 

5.  What  is  the  heat  of  formation  of  sulphuric  acid  from  its  elements 
in  dilute  aqueous  solution,  i.e.     (H2,  S,  O4,  Aq)  =  ?       Ans.  210,700  cal. 


CHAPTER  XVIII. 

NITROGEN   AND   THE  ATMOSPHERE.— THE   HELIUM   FAMILY- 
NITROGEN,   N. 
At.  Wt.  14.01.  Mol.  Wt.  28.02. 

Nitrogen  is  a  colorless,  odorless  and  tasteless  gas.  At  ordinary  tem- 
peratures it  is  chemically  inert. 

Preparation  and  Properties  of  Nitrogen. 
Experiment  I. — Preparation  of  Nitrogen. 

(a)  Assemble  the  parts  of  the  oxygen  or  hydrogen  generator.     Place 
about  5  grams  of  sodium  or  potassium  nitrite,  NaNO2,  KNO2,  and  2  grams 
of  ammonium  chloride,  NH4C1,  in  the  generating  test  tube  or  flask;  add 
10-15  cm. ^  of  water;  insert  stopper  provided  with  delivery  tube;  clamp 
the  tube  to  ring-stand;  heat  gently  by  waving  the  flame  under  the  tube. 
Be  careful  not  to  overheat  the  mixture.     Fill  three  bottles  with  the  gas  by 
displacement  of  water.     Proceed  with  Exp.  II. 

(1)  NaN02  +  NH4C1    ->  NH4NO,  +  NaCl 

(2)  NH4N02  —  N2  +  2H,0. 

(b)  Alternative  method.     The  gas  may  be  prepared  by  heating  ammo- 
nium nitrite,  NH4NO2,  which  decomposes  into  water  and  nitrogen. 

Experiment  II. — Properties  of  Nitrogen. 

Has  the  gas  color,  odor,  taste?  Is  it  inflammable?  Will  it  support 
combustion?  Does  it  unite  readily  with  other  elements?  Prove  its 
inertness  in  this  respect.  Is  the  gas  heavier  or  lighter  than  air? 

Experiment  III. — (Quant.)  Determination  of  the  Weight  of  a  Liter  of 
Nitrogen. 

Provide  a  round-bottomed  250  cm.s  flask  (Fig.  31)  with  a  rubber  cork 
through  which  passes  a  piece  of  glass  tubing  about  8  cm.  in  length;  attach  a 
piece  of  rubber  tubing  about  5  or  6  cm.  long  to  the  outer  end  of  the  glass  tube 
and  wire  it  firmly.  Pour  about  30  cm.3  of  water  into  the  flask,  then  firmly 
press  the  stopper  into  place;  make  a  mark  on  the  neck  of  the  flask  at  the 
bottom  of  the  stopper  so  as  to  be  able  to  determine  the  exact  contents  of 
the  flask  when  the  stopper  is  in  place.  Boil  the  water  with  a  small  flame 
until  all  the  air  has  been  expelled  from  the  flask.  Allow  the  steam  to 
escape  for  5  or  6  min.,  then  close  the  rubber  tube  with  a  strong  clip,  and 

192 


NITROGEN    AND    THE    ATMOSPHERE. — THE    HELIUM    FAMILY.         193 


quickly  remove  the  flame.  Attach  a  fine  wire  to  the  neck  of  the  flask  by 
which  it  may  be  suspended  during  the  process  of  weighing.  When  the 
flask  has  cooled  to  the  temperature  of  the  laboratory,  wipe  and  carefully 
weigh  it.  Read  the  temperature  and  barometric  pressure  in  the  balance- 
room.  Place  a  beaker  of  water  so  that  the  rubber  siphon  tube  of  an  aspira- 
tor bottle  (Fig.  32)  partially  filled  with  nitrogen  is  made  to  dip  beneath  the 
surface  of  the  water.  Connect  the  flask  with  the  delivery  tube  of  the 
aspirator;  the  clip  is  gradually  opened  allowing  a  slow  stream  of  nitrogen 
to  enter  the  flask.  Now  raise  the  beaker  of  water  so  that  the  water  in  it 
will  be  at  a  higher  level  than  that  in  the  aspirator.  Close  the  clip,  dls- 


FIG.  31. 


FIG.  32. 


connect  from  the  aspirator  delivery  tube.  Open  the  clip  for  an  instant  to 
establish  atmospheric  pressure  in  the  flask,  then  weigh  again.  The  in- 
crease in  weight  represents  the  weight  of  the  nitrogen.  From  the  table 
in  the  appandix  ascertain  the  aqueous  tension  at  the  observed  temperature; 
subtract  this  from  the  barometric  pressure.  Determine  the  volume  of  the 
flask  by  filling  with  water  to  the  mark  placed  on  the  neck  and  weighing  it. 
Calculate  the  weight  of  a  liter  of  nitrogen  at  o°  C.  and  760  mm.  What  is 
the  ratio  of  the  weights  of  equal  volumes  of  hydrogen  and  nitrogen? 
What  is  the  weight  of  the  molar  volume  of  the  gas  ? 

THE    ATMOSPHERE. 

The  atmosphere,  or  air  as  it  is  usually  termed,  is  the  gaseous,  envelope 
surrounding  the  earth.  It  is  a  mixture,  the  components  of  which  may 
be  divided  into  two  classes:  (a)  those  components,  oxygen,  nitrogen,  argon 
and  other  members  of  the  helium  group,  which  are  practically  constant 
in  amount;  (b)  components  which  are  variable  in  amount,  e.g.,  carbon 
dioxide,  water  vapor,  ozone,  dust,  traces  of  ammonia,  minute  quantities 
of  solids  like  ammonium  nitrate  and  ammonium  carbonate,  organic  mat- 
ter, etc. 

The  comparatively  recent  discovery  of  five  gaseous  elements  in  the  air 
has  given  a  new  impetus  to  the  investigation  of  the  composition  of  the 
atmosphere.  It  is  rather  remarkable  that  these  five  elements  constitute 
a  natural  group,  known  as  the  helium  group.  The  student  is  referred  to 
the  well-known  work  by  Sir  William  Ramsay,  "The  Gases  of  the  Atmos- 
phere." 

Components  of  the  Atmosphere. 
13 


I94 


EXPERIMENTAL  CHEMISTRY. 


Experiment  IV. — (Quant.)     Determination  of  the  Amount  of  Oxygen 
in  the  Air.     Nitrogen. 

(a)  Absorption   or  Pyrogallate   Method.     See   Exp.    X.     "  Oxygen." 
Repeat  experiment.     Calculate  the  percentage  of    oxygen  [by  volume 
in  the  air.     Test  the  residual  gas  as  to  odor,  taste,  inflammability.     What 

is  the  color  of  the  gas  ?  Will  it  support  combustion  ? 
Are  its  properties  similar  to  those  of  nitrogen  ?  If  this 
residual  gas  were  pure  nitrogen  instead  of  nitrogen 
mixed  with  traces  of  other  gases,  what  would  be  the  per- 
centage of  nitrogen  in  the  air  ? 

Note. — HempeFs  burette  and  compound  pipette  as 

fc,    .,/ IG"J%  „    N  apparata  in  which  to  measure  and  absorb  the  gas  may 

(Smith  and  Keller.}   , rj          ,  .  ,     -  Al  L    ,     ,  0       T°  ,/ 

be  used  instead  of  that  suggested  above.     See  Hempels 

"Gas  Analysis,"  translated  by  L.  M.  Dennis. 

(b)  Explosion  method.     Substitute  Hempel's  explosion  pipette  (Fig.  33) 
for  the  absorption  pipette  used  in  (a),  and  proceed  as  follows:  To  25  cm.3 
of  air  contained  in  the  burette  add  25  cm.3  of  pure  hydrogen,  then  pass  the 
mixture  into  the  explosion  pipette;  close  the  pipette.    (Instructions.)    Pass 
an  electric  spark  through  the  mixture.     Return  the  residual  gas  to  the 
burette  and  measure.     What  is  the  composition  of  this  gas  ?     How  much 
of  the  contraction  was  due  to  oxygen  ?     Calculate  the  percentage  of  oxygen 
in  the  air. 

Experiment  V. — Removal  of  Oxygen  from  the  Air.     Determination  of 
the  Approximate  Percentage  of  Nitrogen  in  the  Air. 

(a)  Place  a  small  piece  of  yellow  phosphorus  either  in  an  evaporating 
dish  floating  on  the  water  or  in  a  small  crucible  resting  in  a  depression 


FIG.  34. 

in  a  large  cork  which  floats  on  the  water  in  the  pneumatic  trough.  Ignite 
the  phosphorus  and  quickly  place  a  bell-jar  or  an  inverted  beaker  over 
the  crucible  (Fig.  34).  Keep  the  rim  below  the  surface  of  the  water  that  no 
gas  may  escape.  Hold  the  bell-jar  in  this  position  until  the  dense  heavy 


NITROGEN   AND    THE    ATMOSPHERE. — THE   HELIUM    FAMILY.         195 


fumes  of  phosphorus  pentoxide,  P2O5,  disappear.  Note  that  the  water  has 
risen  in  the  jar.  Lower  the  bell-jar  until  the  water  on  the  inside  and  out- 
side have  the  same  level;  mark  this  level  on  the  jar.  Place  a  glass  plate 
over  the  mouth  of  the  bell- jar  and  quickly  invert  it;  test  the  gas  with  a 
burning  match.  Results?  Test  the  water  in  the  trough  with  blue  lit- 
mus paper.  Explain.  How  many  cm.3  of  water  are  required  to  fill  the 
jar  to  the  mark  placed  in  it?  How  many  cm.3  are  re- 
quired to  fill  it?  Calculate  the  percentage  of  nitrogen  in 
the  air.  Of  Oxygen.  Equations? 

(b)  If  a  piece  of  yellow  phosphorus  is  inserted  into  a 
measured  volume  of  air  contained  in  a  graduated  tube 
(Fig.  35)  over  water  or  mercury,  the  phosphorus  will  com- 
bine slowly  with  the  air  forming  P2O5  which  unites  with 
water  to  form  phosphoric  acid.  The  phosphorus  should  be 
allowed  to  remain  in  contact  with  the  enclosed  air  for  24 
to  48  hours.  The  water  will  slowly  rise  in  tube  to  take 
the  place  of  oxygen.  In  determining  the  volume  of  the 
residual  gas  the  usual  corrections  should  be  applied.  The 
gas  may  be  tested  for  nitrogen. 

Experiment  VI. — Preparation  of  "Atmospheric  Nitro- 
gen." 

Fit  each  end  of  a  piece  of  hard  glass  tubing  25-30  cm. 
long,  with  a  cork  through  which  passes  a  short  piece  of 
glass  tubing;  clamp  the  tube  in  a  horizontal  position;  *IG'  35- 
place  a  large  plug  of  copper  turnings  in  the  middle  of  the 
tube;  connect  tube  in  series  with  wash  bottles,  No.  i  and  No.  2,  con- 
taining respectively  a  strong  solution  of  KOH  and  concentrated  H2SO4; 
the  latter  wash  bottle  is  connected  with  an  aspirating  bottle  filled  with 
mercury  or  water.  Heat  the  copper  red-hot;  open  the  stop  cock  on  the 
siphon  of  the  aspirator  and  allow  the  liquid  to  be  siphoned  off  drop  by 
drop  or  in  a  slow  stream.  When  the  system  has  been  swept  free  of  air, 
quickly  substitute  another  aspirator  which  has  been  previously  filled. 
When  three-fourths  of  the  liquid  has  run  out,  close  stop  cock  on  siphon  and 
disconnect  the  aspirator  from  the  wash  bottle  and  attach  a  delivery  tube; 
raise  the  end  of  the  siphon  and  attach  a  small  funnel;  pour  water  into 
the  funnel;  open  stop  cock  on  siphon;  the  water  will  drive  the  gas  out 
through  the  delivery  tube.  Collect  several  bottles  of  the  gas  by  displace- 
ment of  water,  and  apply  the  usual  tests  for  nitrogen.  Examine  the 
copper.  Explain  change  of  color. 

Experiment  VII. — Presence  of  Water  Vapor  in  the  Atmosphere. 

(a)  Place  small  quantities  of  calcium  chloride,  CaCl2,  and  phosphorus 
pentoxide,  P2O5,  on  separate  watch-glasses  and  expose  them  to  the  air. 
Examine  them  after  10  or  15  min.  Results?  Conclusions?  Touch  the 
moist  P^,O5  with  a  piece  of  blue  litmus  paper.  Results?  Explain. 
Equation  ? 


196  EXPERIMENTAL  CHEMISTRY. 

Which  of  the  two  substances  apparently  absorbs  moisture  the  more  rap- 
idly ?  Name  another  familiar  substance  which  will  remove  moisture  from 
the  air. 

(b)  A  quantitative  determination  of  the  amount  of  moisture  in  the  air 
may  be  made  by  drawing  a  given  volume  of  air  through  a  U-tube  filled 
with  glass  beads  and  P2O5  by  means  of  an  aspirator.  The  increase  in 
weight  of  the  P2O5  will  represent  the  amount  of  moisture  absorbed  from  a 
known  volume  of  air.  (Instructions.)  Substances  which  absorb  moist- 
ure are  said  to  be  hygroscopic. 

Experiment  VIII. — Presence  of  Carbon  Dioxide  in  the  Atmosphere. 

(a)  Place  5  cm.3  of  a  clear  solution  of  calcium  hydroxide,  Ca(OH)2, 
or  barium  hydroxide,  Ba(OH)2,  in  a  test  tube.     By  use  of  a  glass  tube 
force  air  (carbon  dioxide)  from  the  lungs  through  the  solution  until  the 
latter  has  a  milk-like  appearance,  due  to  the  formation  of  calcium  car- 
bonate, CaCO3.     Equation? 

Half  fill  another  test  tube  or  a  wash  bottle  with  a  Ca(OH)2  solution 
(lime-water)  and  draw  air  through  the  clear  solution  for  several  minutes  by 
means  of  a  filter  pump  or  an  aspirator.  (If  all  the  apparatus  is  not  avail- 
able expose  the  solution  to  the  air.)  Results?  Conclusions  as  to  the 
presence  of  CO2  in  the  air? 

(b)  Repeat  the  above  experiment,  but  pass  the  air  through  a  strong 
solution  of  NaOH  or  KOH  before  it  is  allowed  to  enter  the  Ca(OH)2 
solution.     Results  ?     Explain. 

(c)  The  student  is  referred  to  works  on  "gas  analysis"  for  methods  of 
making  a  quantitative  determination  of  the  CO2  in  the  air.     Report  one 
method. 

Experiment  IX. — Absorption  of  Oxygen  and  Nitrogen  by  Magnesium 
and  Calcium.  Argon,  etc. 

Prepare  a  mixture  composed  of  equal  parts  of  magnesium  powder  and 
freshly-ignited  calcium  oxide.  Keep  the  mixture  perfectly  dry.  Intro- 
duce a  few  grams  of  this  mixture  into  a  piece  of  combustion  tubing  sealed 
at  one  end  and  about  20  cm.  in  length;  clamp  the  tube  in  a  horizontal 
position;  connect  the  open  end  of  the  tube  with  a  delivery  tube  of  smaller 
bore  and  about  80  cm.  long.  The  delivery  tube  should  be  bent  so  that 
its  free  end  may  be  immersed  in  a  vessel  of  mercury  76  cm.  below  the 
combustion  tubing.  Heat  the  mixture  gently,  then  gradually  increase  the 
temperature  to  that  of  a  powerful  burner.  Continue  to  heat  the  tube 
until  the  mercury  no  longer  rises  in  the  vertical  tube.  What  is  the  resid- 
ual gas  ?  What  is  the  probable  composition  of  the  solid  matter  in  the 
tube? 

Experiment  X. — Table  of  the  Composition  of  the  Air. 

Prepare  a  tabulated  statement  of  the  average  composition  of  the  at- 
mosphere. Before  attempting  to  prepare  the  table,  consult  various  ref- 
erence wrorks.  Give  the  names  and  authors  of  the  books  consulted. 


NITROGEN   AND    THE    ATMOSPHERE. — THE   HELIUM    FAMILY.         197 

Experiment  XI. — (Quant.)  Determination  of  the  Weight  of  a  Liter  of  Air. 

The  weight  of  a  liter  of  air  may  be  determined  by  a  method  similar 
to  that  described  under  "  Nitrogen." 

THE    HELIUM    FAMILY. 

The  members  of  this  natural  group  exhibit  many  close  resemblan- 
ces, e.g.,  remarkable  inertness,  monatomic  nature  of  their  molecules, 
etc.  They  also  manifest  that  gradation  of  properties  which  has  been 
observed  in  the  other  natural  groups  of  elements,  e.g.,  atomic  volumes, 
densities,  refractive  indices,  etc.  At  first  there  seemed  to  be  no  place 
within  the  periodic  system  for  these  elements,  but  it  has  been  pointed 
out  that  they  may  not  unnaturally  occupy  the  positions  of  transition 
elements  between  the  highly  negative  elements,  the  halogens,  and  the 
strongly  positive  elements,  the  alkali  metals.  With  the  exception  of  argon 
whose  atomic  weight  appears  to  be  too  large,  the  atomic  weights  of  the 
elements  naturally  place  them  in  this  position. 

Atomic      Absolute  Critical 
Symbol    Density    weight        critical    pressure 

temp.       (mm.) 

Helium  (sun) He  2.0  4.0    

Neon  (the  new  one) Ne          10 .  o         20.0   

Argon  (the  lazy  one) A  19.9         39.9       155.6°        40,200 

Krypton  (the  hidden  one)  ..  Kr          40.9         81.8       210.5°        41,240 

Xenon  (the  stranger) Xe          64.0       128.0       287.8°       43,500 

"  Critical  data"  from  Travers,  Experimental  Study  of  Gases. 

Ramsay  gives  the  following  table  as  representing  an  estimate  of  the 
quantities  in  which  the  individual  gases  of  the  helium  family  are  present 
in  the  atmosphere.  The  data  was  obtained  by  distillation  of  "  atmos- 
phere" air,  or  argon. 

Helium i  to  2  parts  per  1,000,000  of  air. 

Neon i  to  2  parts  per  100,000  of  air. 

Argon °-937  part    per  100  of  air. 

Krypton i  part    per  1,000,000  of  air. 

Xenon i  part  per  20,000,000  of  air. 

Expressed  in  parts  per  1000: 

Argon 9.37 

Neon o.oi 

Helium o .  ooi 

Krupton o .  ooi 

Xenon o .  00005 

It  is  highly  probable  that  the  helium  quickly  passes  out  from  our  at- 
mosphere into  space,  as  a  gas  of  so  low  a  density  must  have,  according 
to  the  "kinetic  theory"  of  gases,  a  velocity  greater  than  the  "critical 
velocity"  at  the  earth's  surface  which  has  a  calculated  value  of  6.9  miles 
per  second.  "The  quantity  of  helium,"  says  Travers,  "which  is  con- 


198  EXPERIMENTAL  CHEMISTRY. 

stantly  being  given  off  by  mineral  springs  is,  however,  enormous,  so  that 
it  is  probable  that  the  amount  present  in  the  atmosphere  does  not  tend 
to  diminish." 

The  student  is  earnestly  requested  to  examine  the  following  books  in 
connection  with  the  work  of  this  chapter. 

Experimental  Study  of  Gases. — Travers.  Gas  Analysis — Hempel 
Dennis.  Liquefaction  of  Gases. — Hardin.  The  Gases  of  the  Atmos- 
phere.— Ramsay. 

PROBLEMS. 

1.  Dumas,  in  determining  the  composition  of  air  by  passing  it  over 
heated  copper  and  measuring  the  residual  nitrogen,  tabulated  the  follow- 
ing data: 

Weight  of  tube  and  copper  before  experiment 120.00  grams. 

Weight  of  tube  and  copper  after  experiment 121.15  grams. 

Weight  of  globe  exhausted 852 .  oo  grams. 

Weight  of  globe  and  nitrogen 855 . 85  grams. 

Calculate  the  percentage  of  nitrogen  and  oxygen  by  weight.  By 
volume. 

2.  Calculate  the  weight  of  5  1.  of  air.     Of  4.2  1.     Of  9  1. 

3.  A  mixture  of  25  cm.3  of  air  and  50  cm.3  of  hydrogen  was  exploded, 
and  the  residue  measured  60.3  cm.3     What  was  the  percentage  of  oxygen 
in  sample  of  air? 

4.  A  student  introduced  50.6  cm.3  of  air  into  a  Hempel  absorption 
pipette  containing  alkaline  pyrogallol.     After  agitating  the  pipette  for 
20  min.  the  gas  was  returned  to  the  Hempel  burette,  and  found  to  occupy 
40.03  cm.3  when  measured  under  original  conditions.     Which  component 
of  the  atmosphere  was  removed  ?     Calculate  the  percentage  of  this  com- 
ponent by  volume. 

5.  Dumas  and  Boussingault  in  1841  found  12.373  grams  of  nitrogen 
and  3.68  grams  of  oxygen  in  a  sample  of  air.     What  per  cent,  of  each 
component  did  they  find? 

6.  A  U-tube   containing   phosphorus   pentoxide  was  found  to  weigh 
30.6293  grams.     A  volume  of  air  which  weighed  30.4268  grams  was 
passed  through,  after  which  the  weight  of  the  tube  was  found  to  be 
31.0517  grams.     Calculate  the  percentage  of  moisture  present  in  the  air. 

7.  It  has  been  found  that  when  air  dissolves  in  water,  the  ratio  of 
oxygen  to  nitrogen  in   the  dissolved  air  is  no  longer  1:3.71  (approx.) 
tut  i :  ?     The  absorption  coefficients  of  oxygen  and  nitrogen  for  water 
as  given  by  Bunsen  are,  O  =  .0411  and  N  =  .0203,  at  760  mm.     Cal- 
culate the  ratio  of  free  oxygen  to  nitrogen  in  the  water  when  the  total 
atmospheric   pressure   is    760   mm.     Hint. — Recall  Henry's   Law   and 
Dalton's  Law  of  Partial  Pressures. 

8. — What  properties  manifested  by  the  atmosphere  led  men  to  believe 
that  it  was  a  chemical  compound? 

Enumerate  the  reasons  which  lead  us  to  regard  the  air  as  a  mixture. 


CHAPTER  XIX. 

NITROGEN  FAMILY. 

Nitrogen N,  14.01 

Phosphorus P,  31.0 

Arsenic As,  75.0 

Antimony Sb,  120.2 

Bismuth Bi,  208 .  o 

These  five  elements  are  grouped  together  not  only  because  of  their 
atomic  weight  relationships,  but  in  virtue  of  a  similarity  in  chemical 
properties.  Antimony  and  bismuth,  especially  the  latter,  possess  de- 
cidedly metallic  characteristics.  With  the  exception  of  bismuth,  they 
form  poisonous  gaseous  derivatives  with  three  atoms  of  hydrogen.  The 
chief  oxygen  compounds  of  each  are  of  R2O3  and  R2O5  types.  They  form 
generally  unstable  derivatives  by  combining  with  three  atoms  of  a  halogen. 
These  halogen  compounds  are  particularly  unstable  in  the  presence  of 
water.  It  should  be  noted  that  with  increase  of  atomic  weight,  the  prop- 
erties of  the  oxides  of  these  elements  vary  gradually  from  strong  acid 
anhydrides  to  weak  basic  oxides. 

NITROGEN,    N. 

At.  Wt.  14.01     Mol.  Wt.  28  02. 
Preparation  and  Properties  oj  Nitrogen. 

Experiment  I. — Recall  Exps.  I  and  II.  "Nitrogen  and  the  Atmo- 
sphere." 

Hydrogen  derivatives,  NH3,  (NH2)2,  N3H. 

Experiment  II. — Preparation  and  Properties  of  Ammonia  Gas  and  Am- 
monium Hydroxide. 

(a)  Preliminary  Experiment.  Dissolve  about  i  gram  of  ammonium 
chloride,  NH4C1,  in  a  few  cm.3  of  water.  Boil  the  solution.  Do  you 
detect  the  odor  of  ammonia  ?  Test  the  escaping  vapors  with  moistened 
pieces  of  red  litmus  paper  and  turmeric  paper.  Results?  Add  a  few 
drops  of  a  NaOH  solution.  Repeat  previous  tests.  Results?  Hold  a 
glass  rod  which  has  been  moistened  with  strong  HC1  just  above  the 
mouth  of  the  test  tube.  Results?  What  is  the  name  of  the  substance 
formed  by  the  interaction  of  the  ammonia  gas  and  HC1?  Equations? 
Repeat,  using  KOH  instead  of  NaOH.  Other  salts  of  ammonia  may  be 
substituted  for  NH4C1. 

109 


2OO  EXPERIMENTAL  CHEMISTRY. 

(b)  Laboratory  method  for  preparation  of  ammonia  gas.     Support  a 
small  flask  by  use  of  wire  gauze  on  a  ring-stand;  fit  a  stopper  provided 
with  a  right-angled  delivery  tube  to  the  flask;  to  the  delivery  tube  con- 
nect by  means  of  rubber  tubing  a  right-angled  tube  which  is  turned  up- 
ward that  the  gas  may  be  collected  by  upward  displacement  of  the  air. 

Mix  together  in  a  mortar  about  20  grams  of  calcium  oxide,  CaO  (quick- 
lime), or  calcium  hydroxide,  Ca(OH)2,  and  10  grms  of  NH4C1.  Odor? 
Place  the  mixture  in  the  flask  and  heat  very  gently;  lay  a  piece  of  card- 
board with  a  hole  in  it  over  the  mouth  of  a  dry  bottle;  invert  the  bottle 
over  the  delivery  tube.  When  the  bottle  is  filled  with  gas,  test  by  wav- 
ing air  from  the  bottle  to  the  nose  or  by  placing  a  piece  of  red  litmus  paper 
moistened  with  water  in  the  mouth  of  the  bottle;  cover  it  and  place  it 
with  mouth  down  upon  the  desk.  Collect  four  bottles  of  the  gas.  Turn 
the  mouth  of  the  delivery  tube  down,  allowing  it  to  barely  touch  the  sur- 
face of  10  cm.3  of  water  in  a  test  tube.  Is  there  any  evidence  that  the 
gas  is  dissolving  in  the  water  ?  Raise  the  test  tube  cautiously  until  the 
mouth  of  the  delivery  tube  is  about  2  cm.  below  the  surface  of  the  water  ? 
Do  bubbles  of  ammonia  gas  escape  from  the  water?  Explain.  Do  you 
notice  any  change  in  the  temperature  of  the  water  in  the  test  tube  ? 

Lower  the  tube  to  its  original  position  and  continue  to  heat  flask  for 
a  few  minutes,  then  remove  the  aqueous  solution  of  the  gas  and  extin- 
guish the  flame.  Raise  the  flask  from  the  gauze  and  place  it  on  a  piece 
of  dry  cloth  and  allow  it  to  cool.  Explain  the  purpose  of  this  last  opera- 
tion. When  the  flask  cools,  observe  if  there  is  any  evidence  of  water  in 
the  generating  flask.  Explain.  Equations? 

(c)  Properties.     In  view  of  the  method  used  in  collecting  the  gas,  wrhat 
are  your  conclusions  as  to  the  density  of  ammonia  gas  as  compared  with 
that  of  the  air?     Test  the  inflammability  of  the  gas.     Results?     Will 
the  gas  support  combustion?     (The  gas  wrill  burn  in  an  atmosphere  of 
pure  oxygen.     Pass  a  stream  of  pure  oxygen  through  concentrated  solu- 
tion of  ammonia  gas  heated  in  a  50  cm.3  flask.     Apply  a  light  to  the  mouth 
of  the  flask.     Explain.     Equations?)     Test  the  gas  with  moist  turmeric 
paper  and  with  litmus  paper.    Results  ?    Moisten  a  glass  rod  with  concen- 
trated HNO3;  insert  it  into  a  bottle  of  the  gas.     Results?     Equations? 

(d)  Solubility  of  the  gas.     Ammonium  hydroxide.     Place  a  bottle  of 
the  gas  mouth  downward  in  a  vessel  of  water.     Hold  it  in  this  position  for 
4  or  5  minutes.    Results.  Explain.  Examine  the  solution  of  ammonia  pre- 
pared in  (b).    Test  it  with  litmus  paper.     Hold  a  rod  dipped  in  concen- 
trated HC1  over  the  solution.     Results?     What  name  is  applied  to  this 
aqueous  solution  of  ammonia  ?     Indicate  its  formation  by  use  of  an  equa- 
tion. 

(e)  Stability  of  ammonium  hydroxide.     Put  5  cm.3  of  NH4OH  in  an 
evaporating  dish  and  boil  it  gently  for  a  few  minutes.     Note  the  escape- 
of  the  ammonia  gas.     Place  5  cm.3  of  NH4OH  in  a  small  beaker  and 
allow  it  to  stand  exposed  to  the  atmosphere  for  24  to  48  hours.     Is 
the  odor  as  strong  as  before  ?     What  do  you  infer  as  to  the  stability  of 
NH4OH? 


NITROGEN   FAMILY.  2OI 

(/)  Does  NH4OH  manifest  the  properties  of  a  base  or  an  acid  ?  Con- 
trast it  with  the  aqueous  solutions  of  hydrogen  compounds  of  sulphur. 
Are  the  basic  properties  of  NH4OH  due  to  the  ammonia  gas  or  the  hy- 
droxyl  group  formed  by  its  combination  with  water  ?  Recall  the  electrical 
conductivity  of  NH4OH.  Is  NH4OH  a  strong  base?  State  the  reasons 
for  your  answer.  What  specific  name  do  we  apply  to  NH4?  Why? 
Write  the  structural  formula  for  NH3  and  NH4OH. 


N  +  H3  —  NH3  +  1 1,900  cal. 
NH3  +  Aq  — »  NH3,Aq  +  8500  cal. 


Ice  will  melt  in  gaseous  ammonia.     Explain. 

Experiment  III. — Production  of  Ammonia  by  Destructive  Distilla- 
tion of  Nitrogenous  Substances. 

Heat  separately  in  a  test  tube  small  quantities  of  each  of  the  following: 
hair,  wool,  bone,  gelatin,  feathers.  Observe  the  effect  of  heat  upon 
these  substances.  Test  the  action  of  the  fumes  which  are  evolved  upon 
turmeric  paper  or  red  litmus  paper.  Results?  Define  destructive 
distillation  ? 

Experiment  IV. — Preparation  of  Ammonium  Salts. 

(a)  Preparation  of  ammonium  chloride.     Place   5  cm.3   of  NH4OH 
(shelf-reagent)  in  a  beaker;  neutralize  the  base  with  dilute  HC1,  using 
litmus  paper  or  phenolphthalein  as  an  indicator;  evaporate  the  solution 
until  crystals  begin  to  form  on  the  sides  of  the  vessel,  then  pour  the  solution 
into  a  crystallizing  dish  and  observe  the  formation  of   crystals  as  the 
solution  cools.     Write  the  formula  for  the  substances  formed. 

(b)  Preparation  of  ammonium  sulphate.     Proceed  as  in  (a),  substituting 
dilute  H2SO4  for  HCL     Are  these  salts  soluble  in  water?     Are  they 
conductors  of  electricity?     If  so,  indicate  by  equations  the  nature  of  the 
electrolytic  dissociation. 

Experiment  V. — Decomposition  of  Ammonium  Salts.     Dissociation. 

(a)  Place  a  small  quantity  of  NH4C1  in  a  piece  of  hard  glass  tubing 
open  at  both  ends.     Clamp  the  tube  in  a  slightly  inclined  position;  heat 
the  tube  strongly  and  hold  moist  pieces  of  each  kind  of  litmus  paper  in 
each  end  of  the  tube.     Results?     Explain.    Recall  Exp.  III.    "Chemical 
Equilibrium." 

(b)  Repeat  (a),  using  (NH4)2SO4  instead  of  NH4C1. 

Nitrogen  and  the  Halogens. 

Experiment  VI. — Preparation  of  an  Endothermic  Compound.  Nitro- 
gen Tri-iodide. 

To  a  small  quantity  of  powdered  iodine  in  an  evaporating  dish,  add  a 
little  strong  NH4OH.  Frequent  stirring  will  assist  the  digesting  of  the 
iodine  by  the  NH4OH.  Let  the  mixture  stand  for  half  an  hour.  Filter. 


202  EXPERIMENTAL  CHEMISTRY. 

Place  the  black  sediment  upon  several  filter  papers.  Spread  the  wet 
papers  at  a  distance  from  one  another  and  allow  them  to  dry.  The 
dry  black  powder  contains  impure  NI3,  which  is  extremely  explosive. 
When  touched  with  a  feather  it  readily  explodes.  The  shock  produced 
by  the  tread  of  a  fly,  or  falling  dust  particles  sometimes  cause  it  to  ex- 
plode. The  student  should  prepare  only  very  small  quantities  of  the  sub- 
stance. 

(1)  2NH4OH  +  I2->  NH4IO  +  NH4I  +  H2O. 

(2)  3NH4IO->N2H3I3  +  KOH. 

(3)  N2H3I3  —   NH3  +  NI3. 

(N,C13)  =    —  38,100  cal. 


Experiment  VII.  —  (Quant.)  Determination  of  the  Weight  of  a  Liter 
of  Ammonia  Gas. 

Fill  a  250  cm.3  flask  with  the  gas  by  upward  displacement  of  air.  Both 
flask  and  gas  must  be  perfectly  dry;  the  gas  may  be  dried  by  passing  it 
through  two  tubes  filled  respectively  with  small  pieces  of  lime  and  soda- 
lime.  A  loose  plug  of  cotton  should  be  placed  in  the  neck  of  the  flask 
to  prevent  diffusion.  When  the  flask  is  filled,  slowly  withdraw  the  tube 
and  cork  at  once.  Wipe  the  flask  and  place  it  in  the  balance-room.  Read 
the  thermometer  and  the  barometer.  Weigh  the  flask.  Determine  the 
volume  of  the  flask.  Calculate  the  weight  of  this  volume  of  air,  if  i  cm.3 
of  air  at  o°  C.  and  760  mm.  weighs  .001293  gram  subtract  this  weight  from 
the  weight  of  the  flask  when  filled  with  air.  The  remainder  is  the  weight 
of  the  vacuous  flask.  What  is  the  weight  of  one  liter  of  the  gas  at  o°  C., 
760  mm.  ?  Of  the  molar  volume  ?  How  many  times  heavier  is  one  liter 
of  ammonia  gas  than  a  liter  of  hydrogen?  What  is  the  ratio  of  their 
molecular  weights  ? 

Experiment  VIII.  —  (Qual.).  Determination  of  the  Composition  of 
Ammonia  Gas. 

Instructions.  —  Perfectly  dry  ammonia  gas  is  passed  very  slowly 
over  heated  magnesium  powder  placed  in  a  piece  of  combustion  tubing. 
Loose  plugs  of  cotton  are  placed  in  the  end  of  the  drying  tubes  contain- 
ing, respectively,  quicklime  and  soda-lime.  The  gaseous  product  is  col- 
lected over  dilute  H2SO4  placed  in  a  small  glass  dish.  The  greenish  pow- 
der which  is  formed  in  the  combustion  tubing  is  tested  by  pouring  a  por- 
tion of  it  into  a  test  tube  half  filled  with  water.  Test  the  gas  which  is 
liberated.  Is  it  ammonia?  Equation? 

It  is  suggested  that  the  ammonia  gas  be  prepared  by  heating  con- 
centrated NH4OH  over  a  low  flame  in  a  generating  flask  provided 
with  a  thistle  tube.  The  flask  should  rest  upon  a  sand  bath  or  a  sheet  of 
asbestos. 

Caution.  —  Do  not  heat  the  thin  layer  of  magnesium  powder  until  the 
apparatus  is  flooded  with  ammonia  gas. 


NITROGEN   FAMILY.  203 

Experiment  IX. — (Quant.)  L.T.  Composition  of  Ammonia  Gas. 
(Volumetric.) 

A  very  desirable  form  of  the  apparatus  necessary  for  the  performance 
of  this  experiment  may  now  be  secured  from  any  of  the  chemical  labora- 
tory supply  houses.  It  is  usually  listed  with  the  "Hofmann  Lecture 
Apparatus."  A  long  glass  tube,  sealed  at  the  lower  end,  and  provided 
at  the  upper  end  with  a  glass  stop  cock,  which  communicates  above  with 
a  small  funnel-like  chamber,  which  can  be  stoppered,  is  rilled  with  chlo- 
rine over  a  saturated  solution  of  sodium  chloride.  The  tube  is  allowed  to 
stand  mouth  downward  in  the  solution  for  sometime  in  order  to  let  the 
liquid  drain  out  of  it,  after  which  the  stop-cock  is  closed  and  the  tube 
on  removal  from  the  solution  is  placed  in  an  upright  position.  The  small 
chamber  above  the  stop  cock  is  now  nearly  filled  with  concentrated  ammo- 
nium hydroxide  and  the  stopper  inserted.  Slightly  open  the  stop  cock 
and  allow  the  ammonium  hydroxide  to  pass  into  the  lower  tube  drop  by 
drop.  The  evolution  of  considerable  heat  and  usually  a  faint  flash  of 
light,  together  with  the  formation  of  dense  white  fumes,  indicate  to  the 
experimenter  that  chemical  reaction  is  taking  place.  Great  care  must  be 
exercised  to  prevent  the  escape  of  any  gas  or  the  entrance  of  any  air  into 
the  tube  when  the  stop  cock  is  opened.  After  nearly  all  of  the  ammonium 
hydroxide  has  been  allowed  to  pass  into  the  tube,  fill  the  funnel  as  before 
and  permit  the  liquid  to  pass  gradually  into  the  tube.  Close  the  stop 
cock.  Fill  a  large  beaker  and  the  funnel  writh  dilute  sulphuric  acid;  fit 
a  tube  bent  twice  at  right  angles  into  a  cork;  fill  the  bent  tube  with  the  acid 
solution  and  fit  the  cork  to  the  funnel;  place  the  other  end  of  this  bent  tube 
in  the  beaker  of  acid.  The  long  tube  should  now  be  immersed  in  a  tall 
cylinder  filled  with  water  at  the  temperature  of  the  laboratory.  Open 
the  stop  cock  and  let  the  acid  drain  slowly  into  the  lower  tube;  if  the 
operation  has  been  successful,  the  acid  will  flow  into  the  tube  until  it  is 
two-thirds  full.  What  is  the  residual  gas  ?  Test  it  with  a  lighted  match. 
What  are  your  conclusions  in  regard  to  the  volumetric  relations  of  the 
two  constituents  of  ammonia  gas  ?  Write  the  equations  for  the  reactions 
involved  in  the  above  experiment. 

NH3,Aq  +  3Cl-*3HCl,Aq  +  N  +  97,300  cal. 

20,600  cal.  3  X  39, 300  cal. 

Experiment  X. — (Quant.)  L.T.  Volumetric  Composition  of  Ammonia 
Gas.  Chemical  Equilibrium. 

(a)  Collect  about  20  cm.3  of  pure  dry  ammonia  gas  in  a  perfectly 
dry  eudiometer  by  displacement  of  mercury.  Allow  the  apparatus  to 
remain  undisturbed  for  10  or  15  min.  The  eudiometer  should  be  clamped 
in  a  vertical  position  in  the  trough  of  mercury  (Fig.  36).  Take  the  read- 
ing of  the  thermometer  suspended  in  the  mercury,  and  record  it  together 
with  the  barometric  reading  and  the  height  of  the  mercury  in  the 
eudiometer.  Reduce  the  volume  of  gas  to  standard  conditions. 


204 


EXPERIMENTAL  CHEMISTRY. 


Connect  the  platinum  wires  of  the  eudiometer  with  an  induction  coil 
and  pass  sparks  through  the  gas  for  30  min.,  or  until  the  volume  of  the 
gas  is  practically  constant.  Allow  the  apparatus  to  remain  undisturbed 
for  10  or  15  min.  then  tabulate  data  as  before.  Reduce  this  final  volume 
to  standard  conditions.  State  the  ratio  of  the  original  and  final  volumes. 
Explain,  indicating  the  involved  reactions  by  means  of  equations,  (b)  or  (c) 
may  be  performed  at  the  discretion  of  the  instructor. 

(b)  Continuation  of  (a).  The  relative  volume  of  hydrogen  and  nitro- 
gen may  be  determined  by  admitting  sufficient  dry  oxygen  to  combine 
with  all  of  the  hydrogen  and  exploding  the  mixture.  The  remaining 


volume,  of  course,  is  nitrogen,  ignoring  aqueous  tension.  It  is  customary 
to  admit  5-10  cm.3  of  oxygen  in  excess  of  that  required  to  combine  with 
the  hydrogen. 

(c)  Chemical  equilibrium.  Continuation  of  (a).  If  a  little  sulphuric 
acid  is  admitted  above  the  mercury  in  the  eudiometer  and  the  mixture 
of  gases  "sparked,"  the  action  will  be  reversed  and  the  volume  of  the 
mixture  will  gradually  decrease  until  all  of  the  gas  disappears.  Explain. 
Interpret  the  following  equations: 


(2%-5%)2NH3<=»N2 
2  NH3  +±  N2 


(95%  - 


,  -*(NH4)2S04. 


Oxygen  Derivatives. 


NITROGEN    FAMILY.  205 

Experiment  XI. — Preparation  and  Properties  of  the  Oxides  of  Nitrogen.. 

(a)  Nitrous  oxide  (nitrogen  monoxide),  N2O.     Place  about  10  grams 
of  ammonium  nitrate,  NH4NO3,  in  the  large  test  tube  used  for  generating 
oxygen;  provide  the  tube  with  a  cork  which  is  fitted  with  a  delivery  tube. 
Heat  the  test  tube  gently.     Collect  three  bottles  of  the  gas  over  warm 
water  by  displacement  of  water.     Observe  the  color  and  odor  of  the  gas. 
Is  it  inflammable  ?     Ascertain  if  it  will  support  combustion  by  lowering  a 
glowing  splint  into  a  bottle  of  the  gas.     Results  ?     What  other  gas  does 
it  resemble  in  its  properties  ?     How  can  they  be  distinguished  ?     Explain 
the  action  of  nitrous  oxide,  N2O,  in  supporting  combustion.     Equations  ? 

(1)  (N2,0)     =— 17,700  cal. 

(2)  (NaO,O)  =  —24,800  cal. 

In  view  of  the  above  experiments  and  equation,  would  you  conclude 
that  N2O  is  stable  or  unstable  ?  Is  it  an  endothermic  or  exothermic  com- 
pound? Indicate  the  nature  of  the  decomposition  of  NH4NO3  by 
means  of  an  equation. 

Note. — Equation  (2)  reveals  the  reason  why  N2O  is  not  reactive;  i.e., 
does  does  not  combine  with  oxygen  to  form  the  higher  oxides. 

(b)  Nitric  oxide  (nitrogen  dioxide),  N2O2  and  NO.     Place  about  15 
grams  of  copper  turnings  in  a  generating  flask  provided  with  a  thistle 
tube  and  a  delivery  tube.     Add   15  cm.3   of  dilute  HNO3,  then  add 
strong  HNO3  until  there  is  a  brisk  evolution  of  gas.     The  addition  of  a 
little  cold  water  will  diminish  the  speed  of  the  reaction.    When  the  brown- 
ish fumes  have  been  swept  out  of  the  apparatus,  fill  four  bottles  with  the 
gas  by  water  displacement.  What  is  the  color  of  the  gas  ?     Lift  one  of 
the  bottles  from  the  water  so  as  to  admit  air.     What  takes  place  ?     What 
is  formed  ?     Is  the  gas  inflammable  ?     Will  it  support  combustion  ?     Test 
by  adding  a  few  drops  of  CS2  to  a  bottle  filled  with  NO;  shake  well,  then 
bring  a  flame  near  the  mouth  of  the  vessel.     Results  ?     Explain.     Would 
you  infer  that  NO  is  stable  or  unstable  ?    Pass  the  gas  into  warm  concen- 
trated HNO3.     Notice  the  formation  of  N2O3  and  N2O4  which  impart 
a    brown,   green  or  blue  color  to   the   solution.     Equation?     Repeat, 
using  a  dilute  solution  of  potassium  permanganate,  KMnO4,  acidified 
with  dilute   H2SO4.     Reoults?     Equations?     Does  the  gas   act  as   an 
oxidizing  or  a  reducing  agent?     Pass  a  current  of    NO  into  a  strong 
solution  of  ferrous  sulphate,  FeSO4.  Results?  Equation? 

(1)  (N,O)  =  —  21,500  cal. 

(2)  (N2,04)=  — 2,700  cal. 

(3)  (2NO,02)=  40,500  cal. 

(4)  (N0,0)    -  13,450  cal. 

(5)  (2NO2,O)  =   2,800  cal. 

After  inspecting  the  above  equations,  what  are  your  conclusions  as 


2O6  EXPERIMENTAL  CHEMISTRY. 

regards  the  readiness  with  which  NO  combines  with  oxygen,  and  con- 
versely, the  ability  of  these  higher  oxides  to  oxidize  ? 

(c)  Nitrogen  trioxide,  N2O3. 

Place  about  one  gram  of  arsenious  acid  (arsenious  oxide),  As2O3,  in  a 
test  tube;  add  8  cm. 3  of  strong  nitric  acid  (1.30-1.35  sp.  gr.)  to  cover 
the  solid,  then  heat  gently.  Note  the  color  and  odor  of  the  evolved  gas, 
N2O3,  as  it  is  conducted  into  a  test  tube  half  filled  with  H2O  Is  the 
gas  soluble?  Test  the  action  of  the  solution  on  litmus  paper. 

Note. — It  is  likely  that  a  small  quantity  of  nitrogen  tetroxide,  N2O4,  is 
produced  simultaneously.  Save  the  solution  for  use  in  Exp.  XVI. 

N2O3  is  the  anhydride  of  what  acid?  See  nitrous  acid.  Indicate  by 
equations  the  reasons  for  you  answer. 

(N2,O3)  =   — 23,000  cal. 
N203-N02  +  NO. 

(d)  Nitrogen  tetroxide  (nitrogen  peroxide),  N2O4  and  NO2. 

Heat  about  8  grams  of  dry  lead  nitrate  in  a  test  tube.  Note  the  color 
and  odor  of  the  gas  evolved.  The  vapor  may  be  condensed  in  a  U-tube 
surrounded  by  a  freezing  mixture,  or  a  portion  of  it  may  be  conducted 
into  a  concentrated  solution  of  NaOH  contained  in  a  test  tube.  If  all 
of  the  gas  is  not  absorbed  by  the  solution,  test  the  escaping  gas  for  oxy- 
gen. Results  ?  What  is  the  name  of  the  gas  ?  Examine  the  substance 
which  remains  in  the  ignition  tube.  What  is  it?  What  is  the  name 
and  formula  of  the  evolved  gas?  Write  equations  for  all  reactions. 
Keep  the  NaOH  solution  for  Exp.  XVII.  Label  it. 

(N,02)  =    -8, 1 25  cal. 

(e)  Nitrogen  pentoxide  (nitric  anhydride),  N2O5,  does   not   exist   in 
the  free  state.     It  is  usually  prepared  by  distilling  a  mixture  of  phosphorus 
pentoxide,  P2O5,   and  nitric  acid.     It  forms  colorless,  rhombic  prisms 
which  are  so  unstable  that  they  explode  violently  when  heated  quickly. 
They  dissolve  readily  in  water  with  the  disengagement  of  considerable 
heat,  forming  nitric  acid,  HNO3. 

2HN03  +  P205  —  N205  +  2HP03. 
(N2,O5  —  gas)  =  — 12, ooo  cal. 
N2O5  +  H2O  —  2HNO3. 

The  series  of  oxides  of  nitrogen  are  an  excellent  illustration  of  what  law 
of  combination?  State  the  law.  Which  of  the  oxides  are  colored? 
Colorless  ? 

(H,N,O3 — liq.)  =  41, 500  cal.;  (H,N,O3 — Aq.)  =  49,000  cal. 

From  the  foregoing  equations  it  is  readily  seen  that  all  of  the  oxides  of 
nitrogen  are  endothermic  compounds,  i.e.,  the  heat  of  formation  from  the 
elements  is  negative.  It  is  obvious  then,  that  the  nitrogen  oxides  cannot 


NITROGEN   FAMILY.  207 

be  prepared  from  the  elements  without  the  addition  of  energy.  (On 
account  of  this  energy  relation,  the  atmosphere  is  able  to  preserve  its 
integrity  as  a  mixture  of  nitrogen  and  oxygen  gases,  and  this  regardless 
of  the  heat  generated  by  frequent  electrical  discharges  in  the  atmos- 
phere.) These  endothermic  oxides  are  unstable,  and  even  explosive. 
Berthelot  exploded  nitric  oxide  by  inflaming  fulminating  mercury. 

An  interesting  bit  of  theory  is  associated  with  the  thermal  relations  of 
nitrogen  and  oxygen.  It  has  been  ascertained  that  the  heat  of  combus- 
tion of  carbon  or  phosphorus  in  nitric  oxide  (NO)  is  about  21,500  calo- 
ries larger  than  when  the  combustion  takes  place  in  oxygen.  This  is 
usually  explained  upon  the  theory  that  the  energy-content  of  nitric  oxide 
is  greater  than  that  of  molecular  oxygen  (O2).  It  is  evident  that  the 
molecules  of  the  gas  supporting  the  combustion  must  be  split  into  their 
constituent  atoms.  This  operation  will  be  attended  by  the  absorption 
or  disengagement  of  heat.  It  follows  then,  that  less  heat  is  required  to 
separate  NO  into  N  and  O  than  to  separate  O2  into  oxygen  atoms.  This 
is  interpreted  as  a  proof  that  the  molecules  of  free  oxygen,  as  of  other  ele- 
ments, possess  an  atomic  structure.  That  is,  the  negative  heat  of  for- 
mation of  nitric  oxide  (N,O  =  —  21,500  cal.)  indicates  that  the  mutual 
affinity  of  the  nitrogen  and  oxygen  atoms  is  less  than  the  sum  of  the  affin- 
ities of  the  oxygen  atoms  for  themselves  and  the  nitrogen  atoms  for  them- 
selves in  their  respective  molecules. 

By  inspecting  these  same  equations  it  will  be  observed  at  once  that  al- 
though the  oxides  of  nitrogen  are  formed  from  the  elements  with  an  ab- 
sorption of  heat,  the  formation  of  the  higher  oxides  from  nitric  oxide  is 
accompanied  with  a  disengagement  of  heat.  This  accounts  for  the 
readiness  with  which  NO  and  N2O3  tend  to  undergo  oxidation,  and  con- 
versely, the  easy  reduction  of  these  higher  oxides  to  nitric  oxide  accounts 
for  their  marked  oxidizing  properties.  (It  will  be  recalled  that  these  ox- 
ides may  act  as  oxidizing  or  reducing  agents.  See  Exp.  XL)  On  the 
other  hand,  nitrous  oxide  (N2O)  unites  with  oxygen  to  form  the  higher 
oxide,  nitric  oxide  (N2O,O),  with  an  absorption  of  heat — therefore  its 
slight  reactivity  with  oxygen.  Its  oxidizing  properties  seem  to  be  due 
wholly  to  its  instability. 

In  concluding  this  note,  the  student  is  reminded  that  the  heat  of  for- 
mation of  compounds  from  the  elements,  as  used  in  this  book,  is  really 
the  heat  of  reaction  (see  note  on  Energetics  of  Chemistry),  and  may  be 
regarded  not  as  an  absolute  but  as  a  relative  measure  of  the  chemical 
energy  of  the  elements  of  which  they  are  composed.  To  illustrate  by 
another  example:  Hydrogen  and  oxygen  unite  to  form  water  as  indi- 
cated by  the  following  equation: 

2H2  +  O2—  2H2O  — gas  +  2  X  57,061  cal. 

Interpreting  the  equation  in  the  light  of  our  theories,  a  definite  quantity 
of  heat  will  be  required  to  split  the  molecules  of  each  gas  into  its  constit- 
uent atoms;  therefore,  the  quantity  of  heat  which  it  is  possible  to  measure 
by  calorimetric  processes  will  be  the  true  heat  of  union  (heat  of  formation) 


208 


EXPERIMENTAL  CHEMISTRY. 


less  this  amount.  This  observed  heat,  which  is  really  the  heat  of  reaction, 
merely  indicates  that  the  mutual  affinity  of  hydrogen  and  oxygen  is  greater 
than  the  sum  of  the  affinities  of  the  hydrogen  atoms  for  themselves  and 
the  oxygen  atoms  for  themselves  in  their  respective  molecules,  by  an 
amount,  the  thermal  equivalent  of  which  is  equal  to  57,061  calories. 
All  affinity  calculations  derived  from  thermal  data  by  such  processes 
must  be  purely  relative. 

Note. — The  heat  of  dissociation  of  the  hydrogen  molecule  into  atoms, 
H2— »(H,H),  has  a  recorded  value  of  128,000  cal.  (Richter.) 

Experiment  XII. — Preparation  of  Nitric  Acid. 

Place  ten  grams  of  pulverized  sodium  nitrate,  NaNO3  in  a  small  tubu- 
lated retort  (Fig.  37)  or  a  distilling  flask;  support  the  retort  on  a  wire  gauze 
in  such  a  manner  that  it  can  be  heated  conveniently.  Introduce  the  end  of 


FIG.  37. 

the  retort  into  a  receiver  (a  dry  test  tube  will  do)  which  is  immersed  in  cold 
water.  Add  20  cm.3  of  strong  H2SO4  by  aid  of  a  funnel  that  no  acid  may 
enter  the  neck  of  the  retort.  Insert  the  stopper  in  the  retort.  Is  there 
any  evidence  of  chemical  action  taking  place?  Heat  the  retort  gently 
and  observe  the  gradual  accumulation  of  liquid  in  the  test  tube.  What 
is  its  color  and  odor?  Action  on  litmus  paper?  Allow  a  few  drops  of 
the  liquid  product  to  run  down  the  inside  of  a  test  tube  containing 
a  solution  of  ferrous  sulphate,  FeSO4.  Results?  Equations?  Repeat 
above  tests  using  the  shelf-reagent,  HNO3.  Results?  What  is  your 


NITROGEN    FAMILY.  2OQ 

conclusion  as  to  the  identity  of  the  liquid  in  the  receiver?  What  is  the 
substance  in  the  retort?  Write  the  equation  for  the  interaction  of  Na- 
NO3  and  H2SO4.  Which  is  the  more  volatile,  H2SO4  or  HNO3? 

Assuming  that  the  reaction  is  reversible,  what  is  the  effect  of  heat 
upon  the  equilibrium  ? 

Experiment  XIII. — Properties  of  Nitric  Acid. 

(a)  Action  of  nitric  acid  on  bases.  To  a  beaker  of  water  add  a  few 
drops  of  NaOH  or  KOH;  test  with  red  litmus  paper.  Now  add  a  dilute 
solution  of  HNO3,  drop  by  drop,  until  neither  blue  nor  red  litmus  paper 
is  altered  in  color  when  dipped  into  the  solution.  Taste  a  drop  of  the 
solution.  What  has  been  formed  by  the  interaction  of  NaOH  and  HNO3  ? 
Write  the  "ionic"  equation  for  the  interaction. 

(6)  Oxidizing  action.  Test  its  action  on  indigo.  Equation  ?  Recall 
the  action  of  concentrated  HNO3  on  iodine.  Equation  ?  Try  the  action 
of  strong  boiling  HNO3  on  powdered  sulphur.  Is  there  any  evidence  of 
chemical  action  ?  Dilute  the  solution,  and  test  for  a  sulphate.  Results? 
Equations?  To  a  solution  of  ferrous  chloride  or  sulphate  add  a  little 
NH4OH.  Result?  Allow  contents  of  tube  to  remain  exposed  to  the  air. 
Results?  Repeat  using  ferric  chloride.  Results?  To  a  solution  of 
ferrous  sulphate  add  a  few  drops  of  concentrated  HNO3;  heat  gently  and 
then  add  NH4OH.  Results?  Explain  the  action  of  HNO3  upon  the 
"  ous  "  salt.  Equations  ?  What  property  of  HNO3  which  is  particularly 
characteristic  of  it,  is  emphasized  in  these  reactions  ? 

(c)  Reduction  of  HNO3.     Formation  of  ammonia.     Place  a  few  grams 
of  zinc  dust  in  a  test  tube,  and  add  a  solution  of  KOH  (i  of  KOH  to  20 
of  H2O).     Warm  the  mixture  gently,  and  test  the  evolved  gas.     What  is 
it  ?     Add  a  few  drops  of  dilute  HNO3.     Note  the  odor  of  the  gas  which  is 
now  given   off.     Confirm   your  conclusions  by  using  two   other   tests. 
Results?     Explain.     Equations? 

(d)  Try  the  action  of  both  dilute  and  concentrated  HNO3  upon  each 
of  the  following  metals:  zinc,  iron,  copper,  lead  and  tin.     Note  in  each 
case  the  kind  of  gas  evolved.     Be  on  the  alert  for  hydrogen,   oxides 
of  nitrogen  and  ammonia.     Devise  a  method  for  showing  whether  the 
product  of  the  reaction  remaining  in  the  test  tube  in  each  case  is  a  nitrate 
or  not.     Apply  it.     Equations?     Compare  the  interaction  of  HNO3  and 
the  metals  with  its  action  toward  the  non-metals.     Refer  to  lecture  notes 
and  reference  texts  for  assistance  in  balancing  equations. 

(e)  Action  on  nitrogenous  substances.     Dip  a  quill,  a  piece  of  wool 
or  pieces  of  white  feather,  in  concentrated  nitric  acid.     The  yellow  color 
is  due  to  xanthoproteic  acid,  formed  by  the  interaction.     This  is  regarded 
as  a  test  for  nitrogen. 

(/)  Determine  the  relative  electrical  conductivity  of  a  5N  solution  of 
HNO3.  Results?  What  is  its  percentage  of  dissociation  in  a  iN 
solution?  (Refer  to  tables.)  Would  you  conclude  that  the  acid  is  very 
active  or  relatively  inactive?  Why? 

Enumerate  several  chemical  properties  of  HNO3  which  are  especially 
14 


210  EXPERIMENTAL  CHEMISTRY. 

characteristic.     Structural  formula  for  nitric  acid?     Anhydride  of  the 
acid? 

Note. — The  complexities  of  oxidation  by  nitric  acid  are  of  such  nature 
that  the  student  usually  requires  some  assistance  in  mastering  the  essential 
principles  of  the  reactions  involved.  The  principles  indicated  in  the 
following  equations  should  be  thoroughly  understood.  The  gaseous 
products  resulting  from  the  interaction  of  HNO3  and  the  metals  and  non- 
metals  are  determined  largely  by  the  concentration  of  the  acid. 

For  dilute  nitric  acid: 

2HNO3—  2NO  +  H9O  +  (3  O) 
S  +  (3  O)  +  H20  —  H2S04. 


2HNO3  +  S  ->  H2SO4  +  2NO. 

For  concentrated  nitric  acid  : 

3(2HNO3)  —  2NO2  +'  H2O  +  (O) 
S  +  3(0)  +  H20  —  H2S04 

6HNO3  +  S  —  H2SO4  +  2H2O  +  6NO2. 

Very  dilute  nitric  acid: 

Zn  +  2HNO3  —  Zn(NO3)2  +  (2H) 
HNO3  +  68(H)  —  NH3  +  3H2O 
or  NH3  +  HN03  —  NH4NO3. 

Dilute  nitric  acid: 

2HN03->H20  +  2NO  +  (30) 
3Cu  +  (30)->3CuO 
3CuO  +  6HNO3  —  3Cu(NO3)2  +  3H2O 

3Cu  +  8HNO3-»  Cu(NO3)2  +  4H2O  +  2NO. 

Concentrated  nitric  acid: 

2HNO3    —  H2O  +  2NO2  +  (O) 
Cu  +  (O)  —  CuO 
CuO  -t-  2HNO3     ->  Cu(NO3)2  +  H2O 

Cu  +  4HNO3  ->  Cu(NO3)2  +  2H2O  +  2NO2. 

The  oxidation  of  "ous"  salts  to  "ic"  salts  by  means  of  nitric  acid  may 
be  represented  by  equations  as  follows: 


3Fe(N03)2  +  4HN03  —  3Fe(N03)3  +  2H2O  +  NO, 
or  2Fe(NO3)2  +  4HNO3  —  2Fe(NO3)3  +  2H2O  +  2NO2. 


NITROGEN    FAMILY.  211 

According  to  the  "ion  theory:" 


*HNO3<=±NO'2  +  OH' 

HNO3  +  H2O  <=±  NO'  '  '  +  3OH' 

3Fe  •  •  +  6NO'3  +  NjD^+^OH'  +  3H"  +  3NO'3  =  3Fe  •  •  • 

9NO'3  +  3H20  +  NO. 

or  2Fe  '  '  +  4NO'3  +  2NO2  +  OH'  -f  2H'  +  2NO'3  =  2Fe  •  •  •  + 

6NO'3  +  2H2O  ^~ 


It  is  obvious  that  oxidation  is  a  process  whereby  the  number  oj  positive 
charges  are  increased,  or  the  negative  are  decreased.  Reduction  consists  oj 
increase  oj  negative  charges  and  a  decrease  oj  the  positive. 

Experiment  XIV.  —  Preparation  and  Properties  of  Aqua  Regia.  Oxida- 
tion of  Hydrochloric  Acid  by  Nitric  Acid. 

To  one  volume  of  concentrated  HNO3  add  three  volumes  of  strong  HC1. 
Gently  warm  the  mixture  and  notice  its  appearance  and  odor.  This 
mixture 

HNO3  +  3HC1—  2H2O  +  NOC1  +  Cl2, 

of  acids  is  usually  referred  to  as  aqua  regia  (royal  water),  because  it  will  dis- 
solve the  noble  metals,  gold,  platinum,  etc.,  forming  chlorides.  No  single 
acid  with  the  exception  of  selenic  acid,  will  dissolve  gold.  Place  a  small 
piece  of  gold  leaf  in  each  of  two  test  tubes.  To  one  add  2  cm.  3  of  HNO3, 
and  to  the  other  6  cm.3  of  HC1.  Observe  that  the  gold  leaf  is  not  attacked. 
Mix  the  contents  of  both  test  tubes  and  warm  gently.  Results?  What 
is  formed  ?  Equations  ?  What  is  the  relative  solution  tension  of  gold  ? 
(refer  to  table  of  solution-tensions).  Explain  the  solubility  of  the  "  noble  " 
metals  in  aqua  regia. 

Experiment  XV.—  Effect  of  Heat  on  Nitrates.  Oxidizing  Power  of 
Nitrates. 

(a)  Recall  Exp.  XI  (a).     Equation  ?     The  nature  of  the  decomposition 
is  peculiar  to  ammonium  nitrate. 

(b)  Recall  Exp.  XI  (d).    Equation?  The  result  is  typical  of  the  nitrates 
of  the  heavy  metals. 

(c)  Heat  sodium  nitrate  in  a  hard  glass  test  tube  before  the  flame  of  a 
blast-lamp  if  necessary  to  secure  an  evolution  of  gas.     Test  the  gas  for 
oxygen.     Continue   to   heat  until  gas  is  no  longer  evolved.     What   is 
the  composition  of  the  residue  ?     (Do  not  attempt  to  answer  this  question 
until  you  have  performed  Exp.  XVI.)     The  action  of  sodium  nitrate  is 
characteristic  of  the  nitrates  of  the  alkali  metals  and  the  alkali-earth 
metals.     Keep  the  residue  for  Exp.  XVI. 

Note.  —  A  residue  of  the    same  composition  may  be  prepared  more 
easily  by  melting  the  nitrate  with  a  piece  of  lead  in  a  crucible,  and  keeping 
*  Ostwald,  Grundriss  der  allgemeinen  Chemie. 


212  EXPERIMENTAL  CHEMISTRY. 

the  mixture  stirred  with  an  iron  spatula.     What  is  action  of  the  nitrate 
on  the  lead? 

Experiment  XVI. — Preparation  and  Properties  of  Nitrous  Acid. 

(a)  Nitrous  acid,  HNO2,  is  formed  by  the  interaction  of  a  nitrite,  say 
sodium  nitrite,  NaNO3,  and  H2SO4. 

Add  a  small  quantity  of  dilute  H2SO4  to  a  dilute  solution  of  NaNO2. 
Test  for  HNO2  as  follows:  Add  a  few  drops  of  the  mixture  supposed  to 
contain  HNO2  to  a  solution  of  KI,  then  add  2  cm.3  of  CS2  and  shake 
vigorously.  Free  iodine  proves  presence  of  HNO2.  Equations? 
Prove  that  a  nitrate  will  not  give  the  same  results.  Record  your  data. 

Try  the  action  of  the  HNO2  on  a  dilute  solution  of  KMnO4.  Results? 
Equation  ?  Does  HNO2  act  as  an  oxidizer  or  as  a  reducing  agent  relative 
toKI?  To  KMnO4? 

(b)  Test  the  aqueous  solution  of  N2O3  prepared  in  Exp.  XI  (c)  for  the 
presence  of  HNO2.     Results?     Equations?     What  is  the  anhydride  of 
HNO2? 

(c)  Dissolve  the  residue  from  heating  NaNO3  (Exp.  XV)  in  a  small 
quantity  of  water.     Filter.     Add  dilute  H2SO4  to  the  nitrate  and  test  for 
the  presence  of  HNO2.     Results  ?     Equation  ?    Was  the  residue  a  nitrate 
or  a  nitrite?     What  is  the  effect  of   heat  on  NaNO3?     Record  your 
answer  under  (c),  Exp.  XV. 

What  class  of  salts  are  formed  by  the  interaction  of  nitrous  acid  and 
the  bases?  Write  the  structural  formula  for  HNO2. 

(H,N,O2  Aq.)  -      -  30,700  cal. 

Experiment  XVII. — Tests  for  Nitrates  and  Nitrites. 

(a)  Test  for  nitrates.     Dissolve  a  small  crystal  of  ferrous  sulphate, 
FeSO4,  in  3  or  4  cm.3  of  water  in  a  test  tube.     Add  a  small  quantity  of 
the  solution  of  any  nitrate.     Now  hold  the  tube  in  a  slanting  position  and 
carefully  pour  down  the  side  3  or  4  cm.3  of  strong  H2SO4.     This  should 
be  done  in  such  a  manner  that  the  acid  may  form  a  layer  at  the  bottom. 
Notice  the  brown  ring  which  forms  at  the  boundary  between  the  two 
liquids.     Explain.     Equation  ? 

(b)  Tests  for  nitrites.     Repeat  (#),  substituting  a  nitrite  for  the  nitrate. 
Results?     Can  this  test  be  used  to  detect  a  nitrite  in  the  presence  of  a 
nitrate,  or  vice  versa?     Recall  the  tests  for  HNO2  (nitrite)  used  in  Exp. 
XVI.     These  tests  will  enable  you  to  detect  a  nitrite  in  the  presence  of  a 
nitrate,  but  the  nitrite  must  be  removed  before  (a)  may  be  applied  as  a 
test  for  the  nitrate.     This  is  usually  accomplished  by  boiling  the  solution 
with  NH4C1.     See  Exp.  I  for  the  reaction. 

The  tests  are  usually  carried  out  as  follows:  3  or  4  cm.3  of  the  "un- 
known" solution  are  placed  in  a  test  tube,  acidulated  with  dilute  H2SO4, 
and  tested  for  a  nitrite,  using  the  KI  or  the  KMnO4  test.  If  a  nitrite  is 
absent,  the  solution  is  tested  at  once  for  a  nitrate;  but  if  a  nitrite  is  present, 
a  portion  of  the  "unknown"  solution  is  boiled  with  NH4C1  until  small 


NITROGEN    FAMILY.  213 

portions  of  the  solution  which  are  removed,  give  no  test  for  the  nitrite. 
The  test  for  nitrates  is  then  applied. 

(c)  Prepare  a  solution  containing  a  nitrite  and  a  nitrate  and  proceed 
with  the  analysis  as  per  directions.     Record  all  data. 

(d)  Test  the  NaOH  solution  of   N2O4  prepared  in  Exp.  XI  (d),  for 
nitrates  and  nitrites.     Results?     What  do  you  infer  was  the  action  of 
the  N2O4  on  the  NaOH  solution? 

All  nitrates  and  nitrites  are  soluble  in  water.  The  silver  salt  of  nitrous 
acid  is  sufficiently  difficult  of  solution  to  be  precipitated  on  the  addition  of 
silver  nitrate  to  a  solution  of  a  nitrite  if  the  latter  is  not  too  dilute. 

PHOSPHORUS,    P. 

At.  Wt.  31     Mol.  Wt.  (PX-P4). 

Phosphorus,  like  sulphur,  exists  in  several  allotropic  forms — crystalline 
or  yellow  phosphorus,  and  amorphous  or  red  phosphorus.  The  properties 
of  these  forms  differ  widely  although  they  are  composed  of  the  same 
materials.  However,  their  respective  energy-contents  are  very  different. 
— See  "Note  on  the  Energetics  of  Chemistry." 

(Yellow)  P  — >  (Red)  P  =  27,300  cal. 

The  approximate  thermal  equivalent  of  the  difference  between  the  in- 
trinsic energies  of  the  allotropic  forms  is  27,300  cal. 

Experiment  I. — Properties  of  Phosphorus. 

Caution. — Yellow  phosphorus  must  be  kept  and  cut  under  water.  It 
must  be  handled  with  forceps,  never  with  the  fingers. 

Determine  the  physical  properties  of  the  yellow  and  red  varieties  of 
phosphorus.  Place  a  quantity  of  powdered  boneblack  upon  an  iron 
plate  (tile  or  brick).  Lay  a  piece  of  yellow  phosphorus  about  the  size 
of  a  pea  upon  this  and  heap  the  powder  up  around  the  phosphorus,  but 
leave  the  top  of  the  phosphorus  exposed  to  the  air.  Does  the  phosphorus 
take  fire?  Explain.  Equations?  Repeat  the  experiment  using  red 
phosphorus.  Results  ?  Explain. 

Determine  the  relative  solubility  of  the  two  forms  of  phosphorus  in 
CS2.  Results?  Dissolve  a  very  small  piece  of  yellow  phosphorus  in 
about  i  cm. 3  of  CS2.  Place  a  filter  paper  on  the  ring  of  the  ring-stand; 
pour  the  solution  upon  the  filter  paper.  Let  the  CS2  evaporate  without 
heating.  Result  ?  Equation  ? 

Prove  that  yellow  phosphorus  is  practically  insoluble  in  water  or  alcohol. 

Drop  a  small  piece  of  yellow  phosphorus  into  a  flask  half  filled  with 
water.  Heat  the  flask  until  the  phosphorus  is  melted,  then  pass  a  current 
of  oxygen  through  a  delivery  tube  upon  the  melted  phosphorus  in.  the 
bottom  of  the  flask.  Results  ? 

Put  a  drop  of  liquid  bromine  in  a  dry  bottle;  allow  the  bromine  to 
vaporize.  When  the  bottle  is  filled  with  the  vapor,  throw  a  small  dry  piece 


214  EXPERIMENTAL  CHEMISTRY. 

of  yellow  phosphorus  into  the  bottle  (?).  Equation?  Is  the  product 
stable?  Repeat,  using  red  phosphorus  (?). 

Bring  together  in  a  porcelain  crucible  or  an  evaporating  dish  a  flake 
of  iodine  and  a  very  small  dry  piece  of  yellow  phosphorus  (?).  What 
is  the  source  of  the  light  and  heat  ?  Equation  ?  Is  the  product  stable  ? 
Does  red  phosphorus  yield  similar  results  ? 

Place  a  small  quantity  of  red  phosphorus  in  the  bottom  of  a  test  tube. 
Clamp  the  tube  in  a  horizontal  position,  and  gently  heat  the  end  containing 
the  phosphorus.  What  is  the  yellow  substance  which  collects  on  the 
cold  portion  of  the  tube  ?  Verify  your  conclusion  by  suitable  tests.  Did 
the  red  phosphorus  melt  or  sublime? 

Tabulate  in  vertical  columns  against  one  another  the  corresponding 
properties  of  these  two  modifications  of  phosphorus.  To  what  is  this 
difference  of  behavior  attributed? 

Hydrogen  Derivatives. 

Experiment  II. — (Hood)     Preparation  of  Phosphine. 

(a)  Into  a  beaker  nearly  full  of  water  (a  few  drops  of  HC1  will  increase 
the  speed  of  the  reaction)  drop  a  small  piece  of  calcium  phosphide,  Ca3P2 
(?).     Equation?     Recall    the    action    of    magnesium    nitride,    Mg3N2, 
upon  water. 

Note. — Calcium  phosphide  has  an  irregular  composition,  and  because 
of  this  fact,  a  mixture  of  the  three  hydrides,  PH3,  P2H4,  P4H2  is  obtained. 
Which  of  the  three  inflames  spontaneously  in  the  air  and  gives  to  gaseous 
phosphine  its  spontaneous  inflammability?  By  what  other  names  is 
gaseous  phosphine  known  ? 

(b)  A  small  generating  flask  (100-250  cm.3)  is  fitted  with  a  rubber 
stopper  provided  with  two  right-angled  delively  tubes,  a  and  b  (Fig.  38). 

A  short  piece  of  rubber  tubing,  carrying 
a  pinch  cock  is  attached  to  a.  In  the 
flask,  place  a  strong  solution  of  KOH 
(i  of  KOH  to  2  of  H2O).  The  flask 
should  be  half  filled,  then  place  it  on  a 
piece  of  iron  gauze  supported  on  the  ring 
stand  in  such  manner  that  the  flask  may 
be  heated.  Connect  b  with  a  delivery 
-  tube  bent  so  that  its  lower  end,  which  is 
FIG.  38.  (Smith  and  Keller.)  turned  upward,  dips  beneath  the  surface 

of  water  in  a  large  beaker  or  a  pneumatic 

trough.  The  water  should  have  a  temperature  of  about  20°.  Now  drop 
a  piece  of  phosphorus  about  the  size  of  a  pea  into  the  flask  and  stopper, 
air-tight.  Connect  a  with  a  hydrogen  generator  and  pass  a  current  of 
hydrogen  through  the  apparatus  until  all  of  the  air  is  displaced;  disconnect 
from  the  generator  and  close  pinch  cock  on  a.  (A  current  of  coal  gas  may 
be  used  instead  of  the  hydrogen,  or  a  few  drops  of  ether  may  be  added  to 


NITROGEN    FAMILY.  215 

the  contents  of  the  flask.  The  ether  on  evaporating  will  drive  the  air  out 
of  the  flask.)  Heat  the  flask  gently  and  phosphine  will  be  evolved  (?). 
Equation?  What  is  formed  by  the  burning  of  the  phosphine?  Equation? 

Caution. — Do  not  remove  the  lamp  or  the  tube  from  the  water  until 
the  reaction  has  terminated,  then  lift  tube  from  water  first,  and  remove 
the  lamp.  The  gas  is  poisonous. 

Read  "Note"  in  (a)  if  you  have  not  already  done  so.  Is  there  any 
similarity  between  ammonia  and  phosphine? 

(Yellow)  P  +  3H  —  PH3  +  11,600  cal. 

Phosphorus  and  the  Halogens. — We  have  seen  that  the  halides  of  phos- 
phorus may  be  formed  by  the  direct  union  of  the  elements.  The  halides, 
being  exothermic  compounds,  are  very  much  more  stable  than  the  analo- 
gous compounds  of  nitrogen.  In  contact  with  water  they  are  unstable, 
undergoing  hydrolysis,  with  the  formation  of  a  hydrogen  halide  and  an 
oxacid  of  phosphorus.  The  reactivity  of  the  halides  of  phosphorus  with 
water  is  fully  explained  by  the  large  amount  of  heat  liberated  at  the  time 
of  the  action. 

(P,C13)   =   75,300  cal.  (P,C15)   =   104,900  cal. 

(P,Br3)  =  42,600  cal.  (P,  I3)     =     10,900  cal. 

Experiment  III.  —  Decomposition  of  the  Halides  of  Phosphorus  by 
Water.  Hydrolysis. 

(a)  Preparation    of    phosphorous    acid  by  hydrolysis  of    phosphorus 
trichloride.     Place  a  few  drops  of  phosphorus  trichloride,  PC13,  in  a  test 
tube;  blow  your  breath  across  the  mouth  of  the  test  tube  (?).     Recall 
the  test  for  a  hydrogen  halide.     Now  add  a  small  quantity  of  water  to  the 
oily  liquid  ( ?).     Ascertain  whether  any  thermal  phenomena  accompanies 
the  dissolving  of  the  PC13  or  not  ( ?).     Warm  the  tube,  and  again  blow  your 
breath  across  its  mouth  (?).     What  are  your  inferences?     Equations? 
Evaporate  the  solution  to  dryness  on  a  water-bath  until  all  of  the  HC1 
has  passed  off.     What  is  the  residue  ?     Add  a  little  water  to  the  residue 
and  try  the  effect  of  the  solution  upon  a  AgNO3  solution.     A  black  pre- 
cipitate proves  the  presence  of  phosphorous  acid,  H3PO3. 

(b)  Preparation  of  phosphoric  acid  by  hydrolysis  of  phosphorus  penta- 
chloride. 

By  means  of  a  spatula,  place  a  small  quantity  of  PC15  in  a  test  tube. 
Blow  your  breath  across  the  mouth  of  the  tube  (?).  Add  a  few  cm.3  of 
water.  Note  the  hissing  noise  which  accompanies  the  reaction.  Is  the 
test  tube  warmed  by  the  interaction  of  the  substances  ?  Warm  the  tube 
and  again  blow  your  breath  across  its  mouth  (?).  Test  it  with  litmus 
paper  ( ?).  Boil  until  the  solution  is  free  of  HC1.  What  is  the  remaining 
fluid?  Add  a  few  drops  of  it  to  a  AgNO3  solution  (?).  A  yellow  pre- 
cipitate of  silver  orthophosphate,  Ag3PO4,  proves  the  presence  of  phos- 
phoric acid,  H3PO4.  Equations  ? 


2l6  EXPERIMENTAL  CHEMISTRY. 

Recall  the  methods  used  for  the  preparation  of  pure  hydriodic  and 
hydrobromic  acids. 

Oxygen  Derivatives. 

Experiment  IV. — Preparation  and  Properties  of  Phosphorus  Pentoiide. 
Phosphoric  Acid. 

Burn  a  carefully  dried  piece  of  phosphorus  under  a  bell- jar.  Allow 
the  jar  to  remain  undisturbed  until  the  heavy  white  vapors  have  deposited 
upon  its  sides.  Compare  the  white  powder  with  the  substance  labeled 
"phosphorus  pentoxide,"  P2O5.  Place  a  portion  of  each  in  separate 
test  tubes  containing  a  little  water  (?).  Is  there  any  thermal  evidence  of 
chemical  action?  Test  each  with  litmus  paper  (?).  Compare  their 
actions  toward  a  silver  nitrate  solution  (?).  When  phosphorus  burns 
with  a  free  supply  of  oxygen,  what  compound  is  formed?  Formula? 
When  this  compound  reacts  with  water,  what  is  the  product  ?  Equation  ? 
Name  two  other  oxides  of  phosphorus  and  give  the  formula  for  each. 
What  is  the  anhydride  of  phosphoric  acid  ?  Of  phosphorous  acid  ? 

(P2A)  =  369,800  cal.          (P205,  Aq.)  =  41,600  cal. 

Phosphorus  pentoxide  is  one  of  the  most  effective  drying  agents  known. 
It  has  been  shown  that  a  glass  tube,  four  or  five  inches  in  length,  filled 
with  the  phosphorus  pentoxide,  will  entirely  dry  a  gas  which  is  slowly 
passing  through. 

Experiment  V. — Preparation  and  Properties  of  Phosphoric  Acid. 

Phosphoric  acid  may  be  prepared  by  heating  a  very  small  quantity  of 
red  phosphorus  with  an  excess  of  strong  nitric  acid;  filter  and  remove 
excess  of  water  by  evaporation  on  the  steam  bath.  The  thick  syrup 
which  remains  should  be  dissolved  and  tested  for  phosphoric  acid  writh  a 
silver  nitrate  solution  ( ?). 

What  is  the  formula  for  phosphoric  acid?  How  many  replaceable 
hydrogen  atoms  ?  Is  it  unibasic,  dibasic  or  tribasic  ?  Into  what  ions  does 
it  dissociate  when  placed  in  water.  Is  it  a  "strong"  acid?  (Refer  to 
tables.)  State  the  reasons  for  your  answer.  Write  the  structural  formula 
for  phosphoric  acid.  Give  the  names  and  formulae  of  three  other  acids 
of  phosphorus. 

From  data  given  in  Exp.  IV  calculate  the  "heat  of  formation"  of  phos- 
phoric acid  in  aqueous  solution  from  its  elements. 

Experiment  VI. — Reactions  of  the  Orthophosphates. 

(a)  Dissolve  a  little  disodium  hydrogen  orthophosphate,   Na^RPC^, 
in  water.     Test  a  portion  of  the  solution  with  litmus  paper  ( ?).     Explain. 

(b)  To  this  portion  add  silver  nitrate  in  solution  ( ?).     Filter.     Ascer- 
tain the  solubility  of  the  yellow  precipitate  in  HNO3  ( ?).     In  NH4OH  (  ?). 
Equations  ? 


NITROGEN    FAMILY.  217 

(c)  Test  a  portion  of  the  original  solution  with  ferric  chloride,  FeCl3  ( ?). 
Add  sufficient  HC1  to  dissolve  the  precipitate,  then  add  slowly  NH4OH 
until  the  solution  is  alkaline  ( ?).     Equations? 

(d)  Test  a  solution  of  any  soluble  orthophosphate  with  "  magnesia 
mixture"  (?).     Is  the  precipitate  crystalline?*     Is  the  precipitate  soluble 
in  excess  of  MgSO4  ?     Is^it  reprecipitated  in  the  cold  by  NH4OH  ?     Equa- 
tions ? 

Note. — "Magnesia  mixture"  is  prepared  by  adding  a  little  NH4OH 
and  an  excess  of  NH4C1  to  a  solution  of  magnesium  sulphate,  MgSO4. 

(e)  Add  an  ammonium  molybdate,  (NH4)2MoO4,  solution  to  a  dilute 
solution   of   an   orthophosphate  acidified  with  concentrated  HNO3  ( ?). 
Describe  the  precipitate.     What  is  it  ?     Give  the  formula.     Equations  ? 

Note. — (d)  and  (e)  are  reactions  frequently  employed  in  analytical 
chemistry. 

Experiment  VII. — Preparation  and  Reactions  of  the  Pyrophosphates. 

When  disodium  hydrogen  orthophosphate,  Na2HPO4,  is  heated  to 
about  300°,  each  two  equivalents  of  it  evolve  one  equivalent  of  water  and 
a  neutral  sodium  pyro phosphate  is  formed  as  a  colorless  glassy  mass,  soluble 
in  water. 

Place  a  small  quantity  of  Na2HPO4  in  a  crucible  and  heat  to  redness  ( ?). 
Equations  ?  Dissolve  the  glassy  residue  ( ?)  in  water  and  test  the  solution 
as  per  Exp.  VI.,  (a),  (&),  (c),  (d),  (e).  Tabulate  results.  Equations? 

Test  a  solution  of  H4P2O7  (or  a  pyrophosphate  acidified  with  acetic  acid) 
with  a  clear  solution  of  albumen.  Is  the  latter  coagulated  ? 

Experiment  VIII. — Preparation  and  Reactions  of  the  Metaphosphates. 

(a).  Heat  to  redness  a  little  microcosmic  salt  (sodium  ammonium 
hydrogen  orthophosphate,  NaNH4HPO4),  and  dissolve  the  residue  of 
sodium  metaphosphate  in  water.  Test  the  solution  as  per  Exp..  VI. , 
a,  b,  c,  d,  e.  Tabulate  results.  Equations? 

Test  a  solution  of  HPO3  with  a  clear  solution  of  albumen.  Is  the 
latter  coagulated  ? 

(b)  Bend  the  end  of  a  platinum  wire  into  a  closed  loop;  heat  it  to  red- 
ness, then  plunge  it  into  a  little  microcosmic  salt;  heat  strongly  in  the  outer 
zone  of  the  blow-pipe  flame  until  a  "glassy  bead"  forms  ( ?).  Touch  the 
hot  bead  to  some  powdered  CuO  and  heat  strongly  again.  If  the  opera- 
tion has  been  successful  a  clear  blue-green  bead  will  be  formed.  The 
solubility  of  certain  metallic  oxides  in  the  metaphosphates  and  the  meta- 
borates  is  taken  advantage  of  to  detect  their  presence.  The  color  of  the 
bead  is  characteristic  of  the  oxide. 

*If  the  precipitate  formed  above  is  not  crystalline  add  a  slight  excess  of  HC1  and 
reprecipitate  with  very  dilute  NH4OH. 


2l8  EXPERIMENTAL  CHEMISTRY.    ' 

Experiment  IX. — Tests  to  Distinguish  the  Phosphates. 

Mention  one  reaction  which  will  enable  you  to  distinguish  orthophos- 
phates  from  all  other  phosphates.  Try  it.  Record  method  and  results. 
How  would  you  distinguish  between  a  meta-  and  a  pyrophosphate  ? 
Try  it.  Record  all  data. 

With  the  exception  of  members  of  the  potassium  family,  the  normal 
orthophosphates  and  pyrophosphates  of  all  the  metals  are  insoluble  in 
water. 

ARSENIC,    AS. 

At.  Wt.  75.0     Mol.  Wt.  (As2,  645°-As4,  1,700°). 

Arsenic  is  essentially  an  acid-forming  element  and  is,  therefore,  a  non- 
metal.  In  some  of  its  compounds,  however,  it  functions  apparently 
as  a  metal.  It  cannot  displace  hydrogen  from  dilute  acids.  It  has 
valences  of  three  and  five. 

Experiment  I. — Properties  of  Arsenic. 

(a)  Study  the  physical  properties  of  the  elementary  substance.     Does 
it  possess  a  metallic  appearance  ?     Is  it  crystalline  in  form  ?    Place  a 
piece  of  arsenic  about  half  the  size  of  a  grain  of  wheat  in  a  small  tube  of 
hard  glass  and  heat  to  redness  ( ?).     What  name  is  applied  to  such  change 
of  state  ?     Repeat  using  orpiment,  As2S3,  or  realgar,  As2S2. 

(b)  Heat  a  very  small  piece  of  arsenic  on  charcoal  in  the  oxidizing  flame. 
Results  ?     Note  that  the  characteristic  odor  of  garlic  is  perceptible.     Will 
the  free  element  burn  in  the  air?     If  so,  what  is  formed? 

(c)  To  a  small  quantity  of  arsenic  in  a  test  tube  add  an  excess  of  HNO3. 
Boil  so  long  as  brown  fumes  form,  or  until  the  liquid  does  not  color  on  cool- 
ing.    Save  the  solution  for  use  in  Exp.  V. 

Recall  similar  experiments  with  iodine  and  phosphorus  (?). 

Hydrogen  Derivatives. 

Experiment  II. — Arsine.     Marsh's  Test  for  Arsenic. 

(Hood;  Poison!  Instructions).  Thoroughly  clean  the  hydrogen- 
generator;  place  a  small  quantity  of  chemically  pure  zinc  in  the  flask; 
connect  the  generator  with  a  tube  (U-tube  or  bulb-tube)  filled  with  cal- 
cium chloride,  CaCl2.  Connect  this  latter  tube  with  a  piece  of  hard  glass 
tubing,  r,  about  20  cm.  long,  clamped  in  a  horizontal  position  (Fig.  39).  It 
is  well  to  make  constrictions  in  the  tube  at  intervals  of  7  or  8  cm.  by  gently 
heating  and  drawing  it  out.  Now  add  dilute  HC1  through  the  thistle  tube. 
When  the  air  has  been  displaced  by  hydrogen,  light  the  gas  by  the  "test  tube" 
method.  Observe  the  color  of  the  flame.  By  means  of  a  pair  of  pinchers 
hold  a  porcelain  crucible  lid  in  the  flame.  If  there  is  no  deposit  of  solid 
matter  upon  the  lid,  introduce  through  the  thistle  tube  a  few  drops  of  a 
solution  of  arsenic  trichloride  (arsenic  trioxide,  As2O3,  dissolved  in  HC1). 


NITROGEN    FAMILY. 


219 


Note  the  appearance  of  the  flame  ( ?).  Is  there  a  deposit  upon  the  cold 
porcelain  lid?  Marsh's  test:  Heat  the  horizontal  tube  between  any 
two  constrictions,  with  a  flame.  Results  ?  Locate  the  deposit  in  the  tube 
relative  to  generator  and  the  flame.  Has  the  deposit  upon  the  lid  a  me- 
tallic luster?  Test  the  solubility  of  the  deposit  (arsenic  "spot")  with 
any  soluble  hypochlorite.  Indicate  by  equations  the  chemical  reactions 


FIG.  39. 

involved  in  the  preparation  of  the  gas,  arsine.  Explain  by  use  of  equa- 
tions the  changes  occurring  within  the  flame.  What  other  name  is  some- 
times applied  to  the  gas,  AsH3  ? 

Note. — The  student  may  be  asked  to  determine  whether  or  not  arsenic 
is  present  in  the  green  coloring  matter  used  on  wallpaper  and  shipping- 
labels. 

As  +  H3— *  AsH3  — iijoocal. 

Oxygen  Derivatives. 

Experiment  III. — Arsenic  Trioxide  (Arsenious  Oxide).    Arsenious  Acid. 

(a)  Place  a  little  powdered  arsenic  in  a  hard  glass  tube  open  at  both 
ends  and  about  20  cm.  long.     Clamp  the  tube  in  a  nearly  horizontal  posi- 
tion and  heat.     Notice  the  white  deposit  of  arsenic  trioxide,  As^g,  on 
the  colder  portions  of  the  tube.     Scrape  the  deposit  into  a  small  test  tube 
and  resublime  it.     Does  the  sublimate  show  a  crystalline   structure? 
By  what  other  name  is  arsenic  trioxide  commonly  known  ? 

(As2,  O3)  =  1 54,600  cal. 
(A&j,  O5)  =  2 19,400  cal. 

(b)  Reduction  of  arsenic  trioxide.     Heat  a  pinch  of  As2O3  after  mixing 
it  with  a  little  powdered  wood  charcoal,  in  a  very  narrow  test  tube  (or, 


220  EXPERIMENTAL  CHEMISTRY. 

better,  in  a  drawn-out  glass  tube  having  a  small  bulb  on  the  end).     Re- 
sults ?     Equation  ? 

Note. — It  is  advisable  to  cover  the  mixture  with  a  layer  of  powdered 
charcoal. 

(c)  Preparation  of  arsenic  trichloride.     Boil  a  small  quantity  of  As2- 
O3  with  concentrated  HC1  (?).     Equation?     Add  a  few  cm.3  of  water. 
Preserve  the  solution  for  future  use.     Does  As2O3  in  this  reaction  manifest 
the  properties  of  a  metallic  or  a  non-metallic  oxide,  i.e.,  basic  or  acidic  prop- 
erties ? 

(As,  C13)  =  71,400  cal. 

(d)  Preparation  of  sodium  arsenite,  Na3AsO3.     Boil  a  small  quantity 
of  As2O3  with  a  solution  of  sodium  hydroxide.     Results  ?     Equation  ? 

Does  As2O3  in  this  reaction  manifest  the  properties  of  a  metallic  or  a 
non-metallic  oxide? 

(e)  Formation  of  arsenious  acid,  H3AsO3.     Heat  a  little  As2O3  in  10 
cm.  of  distilled  water  in  a  test  tube.     Filter  if  the  solution  is  not  clear. 
Test   the   action  of   the  nitrate  upon  litmus  paper  (?).     Conclusions? 
As2O3  is  the  anhydride  of  what  acid  ?     Ascertain  the  solubility  of  As2O3  in 
water  by  evaporating  a  portion  of  the  solution  to  dryness  (?).     The 
names  of  the  salts  formed  by  this  acid  have  what  ending?     Is  H3AsO3 
known  in  the  free  state  ?     What  is  the  best  antidote  for  arsenious  oxide  ? 
Write  the  structural  formula  for  arsenious  acid. 

Experiment  IV. — Salts  of  Arsenious  Acid. — Arsenites. 

(a)  To  a  portion  of  a  solution  of  sodium  or  potassium  arsenite  add 
a  little  silver  nitrate  solution  (?).     Filter.     Try  the  effect  of  an  excess 
of  NH4OH  upon  a  portion  of  the  precipitate  ( ?).     Boil  ( ?).     What  is  the 
effect  of  HNO3  upon  the  precipitate?     Equations? 

(b)  Add  a  few  cm.3  of  a  solution  of  copper  sulphate  to  a  portion  of  the 
aqueous  solution  of  the  arsenite  ( ?).     Is  the  precipitate  soluble  in  NaOH  ? 
Heat  the  solution  (?).     Equations? 

Experiment  V. — Arsenic  Acid. — Arsenates. 

Note. — The  student  should  be  on  the  alert  to  trace  any  analogies  of 
crystalline  form,  composition,  solubility,  etc.,  of  the  phosphates  and 
arsenates. 

(a)  Evaporate  the  solution  of  arsenic  in  nitric  acid,  Exp.  I  (c\  to  dry- 
ness  and  redissolve  the  residue  in  20  cm.3  of  hot  water.     The  purpose 
of  this  operation  is  to  remove  any  free  HNO3.     Test  the  solution  with  lit- 
mus paper  ( ?).     Is  the  acid  reaction  due  to  nitric  acid  or  arsenic  acid, 
H3AsO4  ?     Arsenic  acid  may  be  prepared  by  substituting  As2O3  for  the 
metallic  arsenic  and  following  the  foregoing  procedure.     Equations  ? 

(b)  Precipitation  of  silver  orthoarsenate.     Add  a  slight  excess  of  Ag- 
NO3  solution  to  5  cm.3  of  the  arsenic  acid  solution  ( ?).     Filter.     Try  the 
solubility  of  separate  portions  of  the  precipitate  in  NH4OH  (?),  and  in 


NITROGEN    FAMILY.  221 

HNO3(?).     What  is  the  effect  of  boiling  the  NH4OH  solution?     Equa- 
tions ? 

(c)  Precipitation  of  magnesium  ammonium  orthoarsenate.     To  "  mag- 
nesia mixture"  containing  an  excess  of  NH4OH,  add  5  cm.  of  the  solution 
of  H3AsO4  (or  any  soluble  ar senate]  ( ?).     If  the  precipitate  is  not  crystal- 
line dissolve  it  in  the  smallest  quantity  of  HC1,  and  reprecipitate  it  .with 
NH4OH.     Equations? 

(d)  Precipitation  of  ammonium  arseno-molybdate.     Add  an  ammoniun 
molybdate  solution  to  5  cm.3  of  the  H3AsO4  solution  (or  any  soluble  arsen- 
ate)  acidified  with  concentrated  HNO3  ( ?).     If  precipitation  is  apparently 
incomplete,  warm  the  mixture  ( ?).    Equations  ?    Name  and  give  the  form- 
ulae of  four  acids  of  arsenic.     Into  what  ions  does  arsenic  acid  dissociate 
when   placed   in   water?     Is   it   a   "strong"   acid?     (Refer   to   tables.) 
Write  the  structural  formula  for  the  acid. 

Experiment  VI. — (Optional)  Reactions  of  the  Orthoarsenates. 

Exp.  VI.— "Phosphorus, "  may  be  repeated  using  any  soluble  arsenate 
instead  of  Na2HPO4.  Equations? 

Experiment  VII. — Sulphides  of  Arsenic.  Colloidal  Solutions.  Sulpo-salts. 

(a)  Precipitation  and  properties  of  arsenic  trisulphide.     Pass  H2S  into 
a  solution  of  H3AsO3  (or  of  an  arsenite)  (?).     What  name  is  applied  to 
this  kind  of  a  solution  ? 

Pass  H2S  into  the  solution  of  As2O3  in  HC1  prepared  in  Exp.  Ill  (c), 
(or  of  an  arsenite  acidified  with  HC1)  (?).  Filter  and  wash  the  precipi- 
tate. Divide  the  precipitate  into  two  parts.  Boil  one  portion  with  dilute 
HC1.  Is  the  precipitate  apparently  soluble  ?  Place  the  other  portion  of 
the  precipitate  in  an  evaporating  dish  and  add  few  cm.3  of  warm  ammo- 
nium sulphide,  (NH4)2S,  or  warm  ammonium  polysulphide,  (NH4)2SX  ( ?). 
Add  HC1  to  the  solution  thus  formed  and  note  the  effect?  Equations? 

(b)  Precipitation  and  properties  of  arsenic  pentasulphide.     Repeat  (a}, 
substituting  H3AsO4  (or  an  arsenate)  for  the  H3AsO3. 

Experiment  VIII. — Tests  for  Arsenic. 

In  addition  to  Marsh's  test,  the  following  tests  are  used  for  the  detec- 
tion of  arsenic:  Reinsch's,  Bettendorff's,  Gutzeit's,  Fleitmann's.  Con- 
sult works  on  analytical  chemistry  and  make  a  detailed  statement  of  one 
of  the  above  methods.  Test  the  method  before  reporting.  The  method 
may  be  presented  in  the  form  of  a  record  of  your  experiment. 

ANTIMONY,    Sb. 

At.  Wt.  120.  2  Mol.  Wt.  (Sb2-Sb4). 

The  metallic  characteristics  exhibited  by  arsenic  become  more  marked 
with  antimony.  It  is  both  an  acid-forming  and  base-forming  element. 


222  EXPERIMENTAL  CHEMISTRY. 

It  is  a  weak  "metalloid."  It  gives  sets  of  compounds  in  which  it  is  triv- 
alent,  and  others  in  which  it  is  quinquivalent.  It  cannot  displace  hy- 
drogen from  dilute  acids. 

Experiment  I. — Properties  of  Antimony. 

Note. — The  chemical  properties  of  antimony  are  very  similar  to  those 
of  arsenic. 

(a)  Study  the  physical  properties  of  the  elementary  substance.     Does 
it  possess  a  metallic  appearance  ?     Is  it  crystalline  ?     Heat  a  small  piece 
of  antimony  to  redness  in  a  hard  glass  tube  closed  at  one  end  ( ?).     Re- 
peat the  experiment  using  a  tube  open  at  both  ends  (  ?). 

(b)  Heat  a  small  piece  of  antimony  on  charcoal  in  the  oxidizing  flame. 
Results  ?     Will  the  free  element  burn  in  the  air  ?     If  so,  what  is  formed  ? 

(c)  Boil  a  little  powdered  antimony  in  concentrated  HNO3.     Results  ? 
Preserve  for  use  in  Exp.  IV. 

Hydrogen  Derivatives. 

Experiment  II. — Stibine.     Test  for  Antimony. 

Repeat  Marsh's  test,  using  antimony  trichloride,  SbCl3,  in  HC1  solu- 
tion in  place  of  AsCl3.  Distinguish  between  arsine  and  stibine. 

Note. — The  student  may  be  asked  to  prove  experimentally  that  tartar 
emetic  contains  antimony.  What  is  the  formula  of  this  salt? 

Sb  -f  H3  — »  SbH3  —84,500  cal. 
Halides. 

Experiment  III. — Hydrolysis  of  Antimony  Trichloride.     Mass  Action. 

(a)  Recall  Exp.  II.     "Chlorine."     Equation? 

(b)  Place  a  small  crystal  of  antimony  chloride,  SbCl3,  in  a  test  tube. 
Add  10  cm.3  of  water  and  shake  (?).     Test  the  liquid  \vith  litmus  paper 
(?).     Prepare  a  clear  solution  by  adding  concentrated  HC1,  a  drop  at  a 
time,  warming  the  mixture  after  each  drop.     The  solution  contains  what  ? 
To  5  cm.3  of  the  solution  add  a  large  amount  of  water.    Result  ?    Explain. 
What  kind  of  action  is  this  ?     Write  the  equation  for  the  reaction. 

Now  add  concentrated  HC1.  Result  ?  Is  the  action  reversible  ? 
What  influences  the  direction  of  the  reaction  ?  Write  an  equation  which 
will  represent  both  actions.  Referring  to  your  equation,  in  which  direc- 
tion does  the  reaction  go  with  the  greatest  speed  when  an  excess  of  acid  is 
used  ?  When  an  excess  of  water  is  used  ? 

(Sb,Cl3)  =  91, 400  cal. 
Oxygen  Derivatives. 


NITROGEN    FAMILY.  223 

Experiment  IV. — Antimony  Trioxide.     Antimony  Pentoxide. 

Examine  the  residue  which  remained  in  the  tube  after  heating  powdered 
antimony  with  HNO3  (Exp.  I,  (c)).  The  white  insoluble  residue  is  prob- 
ably a  mixture  of  antimony  tri-  (Sb2O3)  and  penta-  (Sb2O5)  oxides  of  anti- 
mony. These  oxides  are  the  anhydrides  of  what  acids?  Filter  the 
solution  and  boil  a  small  portion  of  the  residue  with  a  solution  of  NaOH 
until  a  clear  solution  is  obtained.  Products?  Does  the  oxide  exhibit 
acid  or  basic  properties  in  this  reaction  ?  Equation  ? 

Boil  another  small  portion  of  the  residue  with  strong  HC1.  Products? 
Would  you  infer  from  this  reaction  that  the  oxide  of  antimony  possesses 
basic  properties?  Why?  Evaporate  the  clear  solution  to  small  bulk. 
Does  it  interact  with  water  like  SbCl3?  Save  a  portion  of  the  clear  solu- 
tion for  Exp.  V. 

(Sb2,  0,,  3H20)  =  167,400  cal. 
(Sb2)  O5,  3H2O)  =  228,700  cal. 

Experiment  V. — Sulphides  of  Antimony.     Sulpho-salts. 

Into  a  solution  of  SbCl3  acidified  with  HC1  pass  H2S  (?).  Equation? 
Filter  and  wash  the  precipitate.  Divide  the  precipitate  into  two  parts. 
To  one  portion  add  strong  HC1.  Warm.  Result?  Equation?  Try 
the  action  of  warm  (NH4)SX  upon  the  other  portion  of  the  precipitate  ( ?). 
Equation  ?  To  the  solution  add  HC1  ( ?).  Is  Sb2S3  soluble  in  (NH4)2SX  ? 
What  salt  is  formed?  Equations? 

BISMUTH,    Bi. 

At.  Wt.  208.0         Mol.  Wt.  (Bi-Bia). 

The  metallic  character  of  bismuth  considerably  exceeds  its  metalloidal 
properties.  It  does  not  form  a  hydrogen  derivative,  and  the  oxide  (Bi2O3) 
which  possesses  a  constitution  similar  to  the  acid  forming  As2O3  exhibits 
only  basic  characteristics.  Bismuth  and  its  derivatives  are  usually  con- 
sidered with  the  metals,  but  on  account  of  the  fact  that  it  forms  a  number 
of  compounds  analogous  in  composition  and  properties  to  the  compounds 
of  other  members  of  the  group,  it  is  considered  in  this  order. 

Experiment  I. — Properties  of  Bismuth. 

(a)  Note  the  physical  properties  of  the  elementary  substance   (?). 
Is  it  malleable  or  brittle.     (Test  with  the  pestle.) 

(b)  Heat  a  small  piece  of  bismuth  on  charcoal  in  the  oxidizing  flame. 
Results  ? 

(c)  Mix  a  little  oxide  or  nitrate  of  bismuth  with  a  small  quantity  of 
sodium  carbonate,  Na2CO3.     Heat  the  mixture  on  charcoal  in  the  re- 
ducing flame.     Examine  the  resulting  metallic  globule  (?).     Explain. 
Equation  ? 

(d)  Treat    a    pinch    of    powdered    bismuth    with    HNO3.     Result? 


224  EXPERIMENTAL  CHEMISTRY. 

Products?     Concentrate    the    solution,   cool  and   crystallize.     Save   for 
Exp.   III. 

Experiment  II. — Alloys.     Wood's  Metal. 

Place  a  small  piece  of  Wood's  metal  (an  alloy)  in  a  small  test  tube. 
Support  the  test  tube  in  a  beaker  of  water;  heat  the  water  in  the  beaker. 
Determine  the  temperature  at  which  the  alloy  melts  by  taking  the  tem- 
perature of  the  wrater-bath  (?).  Name  the  components  of  the  alloy  and 
give  their  respective  melting  points. 

Experiment  III. — Hydrolysis  of  Bismuth  Nitrate. 

Dissolve  the  crystals  prepared  in  Exp.  I  (d),  (or  crystals  of  bismuth 
nitrate,  Bi(NO3)3)  by  heat  with  the  addition  of  the  least  possible  amount 
of  nitric  acid.  To  a  fewr  cm.3  of  the  solution  add  a  large  quantity  of 
water.  Result?  Add  a  few  drops  of  dilute  HNO3,  just  enough  to  re- 
move the  cloudy  appearance  (?).  Add  a  large  volume  of  water  (?). 
Explain.  Equations  ?  Is  the  action  reversible  ? 

Experiment  IV. — Preparation  of  Bismuth  Hydroxide.  Dehydration. 
Bismuth  Trioxide. 

(a)  Try  the  effect  of  NH4OH  upon  a  clear  solution  of  Bi(NO3)3. 
Result  ?  What  are  the  products  ?  Equation  ?  Filter  and  wash  the 
precipitate. 

(6)  Remove  the  precipitate  from  the  filter  and  place  it  in  a  porcelain 
crucible.  Ignite.  Note  the  color  of  the  residue  of  bismuth  trioxide 
when  hot;  also  cold  ( ?).  Equation  ?  Define  dehydration. 

(Bi2,03)  =  ?cal. 

Experiment  V. — Preparation  of  Bismuth  Trichloride.     Hydrolysis. 

Pour  a  few  drops  of  dilute  HC1  upon  precipitated  Bi(OH)3  upon  a 
filter  (?).  Collect  the  filtrate  in  a  test  tube  containing  15  to  20  cm.3  of 
water.  Result?  Explain.  Equation?  The  halide  of  what  other  element 
manifests  a  tendency  to  undergo  hydrolysis  ? 

(Bi,Cl3)  =  90,600  cal. 

Experiment  VI. — Bismuth  Trisulphide. 

Pass  H2S  through  a  solution  of  the  chloride  or  nitrate  of  bismuth  con- 
taining just  enough  HC1  to  prevent  hydrolysis.  Result?  Products? 
Equation?  Filter  and  discard  the  filtrate.  Treat  the  precipitate  with 
warm  (NH4)2SX.  Is  the  precipitate  apparently  soluble?  Filter  and 
add  HC1  to  the  filtrate.  Is  the  yellowish-white  substance  free  sulphur 
or  bismuth  trisulphide,  Bi2S3?  Is  Bi2S3  soluble  in  (NH4)2SX?  Equa- 
tions? Compare  results  with  those  in  Exp.  V.  "Antimony"  (?). 

The  following  comparative  table  will  show  that  the  elements  of  the  nitro- 


NITROGEN    FAMILY. 


225 


gen  family  exhibit  the  same  gradation  of  properties  with  increasing  atomic 
weights  as  was  displayed  by  the  elements  of  the  oxygen  family: 


Physical  Properties         Nitrogen 

Atomic  weight 14.01 

State  or  phase Gas 


Color  Colorless 


Phosphorus       Arsenic  Antimony  Bismuth 

31.0                75.0  120.2  205.0 

Solid              Solid  Solid  Solid 

(crystalline)  (crystalline) 

(crystalline)  (crystalline) 

Yellow         Steel-gray  Silver-  Reddish, 

white  Metallic. 


Specific  gravity  
Specific  heat  

.886  (liq 

uid)        1.82 

5-72 

6.7 

9.8 
/ 

*x 

Melting-point 

2  I  7° 

44° 

4CO° 

42^° 

270° 

Boiling-point 

*"      O 

-193° 

TTT" 

287° 

*4-Ow 
800° 

t*j 

1,400°? 

^  /  w 

i,6oo°? 

Chemical  properties 

H-derivatives 

NH3 

PH, 

AsH? 

SbH, 

Heat  of  formation.  . 

1  1,  800 

J.    -t--l-3 

cal.        1  1,  600 

•  1-OJ--L3                   ~  "^^j 

cal.  —  11,700  cal.  —  84,500  cal  

Chemical   properties 

Strong  base          Weak 

base     Neutral 

Neutral 

Halides  

NCI, 

PCI, 

AsCl3 

SbCl3 

BiCl3 

Heat  of  formation  .  . 

•*•  ^  ^-  *3                     ~~  v  ~o 

—  38,100  cal.     75,000  cal 

71,400  cal. 

91,400  cal. 

90,600  cal. 

State  or  phase  

Liquid 

Liquid 

Liquid 

Solid 

Solid 

Stability  

Explosive 

Decomposed 

Decomposed 

Undergoes 

Undergoes 

by  water 

by  water 

hydolysis 

hydrolysis. 

O-derivatives  

N203 

P203 

As203 

Sb2O3 

Bi203 

and  

N20S 

P30s 

As2Os 

Sb2Os 

Bi20s. 

Chemical   properties 

Anhy- 

Anhydrides 

Anhydrides 

Weak  anhy- 

Strongly 

drides 

drides, 

basic. 

basic. 

Acids  

HNO2 

H3PO3 

H3AsO3 

HSbO2 

IVToin  ]]'\r     r\  mt  ^r»rf  ioc 

HNO3 

H3P04 

H3AsO4 

H3SbO4 

^» 

PROBLEMS. 

1.  If  a  liter  of  water  absorbs  800  times  its  volume  of  ammonia  gas, 
how  many  grams  of  ammonia  will  be  absorbed  by  500  cm.3  of  water? 

2.  How  many  grams  do   100  cm.3  of  concentrated  HNO3  weigh? 

3.  How  many  grams   of   concentrated   HNO3  can  be  made  from  10 
grams  of  sodium  nitrate  ? 

4.  How  much  phosphorus  can  be  obtained  from  100  grams  of  bone- 
ash,  containing  75  per  cent,  of  calcium  phosphate,  Ca3(PO4)2?     From 
100  grams  of  pure  apatite,  Ca5F(PO4)3  ? 

5.  Calculate  the  percentage  of  arsenic  in  realgar.     In  orpiment. 

6.  The  specific  gravity  of  the  vapor  of  phosphorus  relative  to  hydrogen 
(H2)  is  119.8  at  800°.     How  many  atoms  in  the  molecule  of  the  elemen- 
tary substance  ? 

7.  Specific    gravity  of  vapor   of    arsenic  at  900°  is  296.6;  at  1740°, 
154.6.     Determine  the  molecular  formula  of  arsenic  at  the  respective 
temperatures. 


'5 


CHAPTER  XX. 
CARBON  FAMILY. 

Carbon C,  12.00 

Silicon Si,  28.4 

(Germanium  .  .  .  Ge,  72.5) 

(Tin Sn,  119.0) 

(Lead Pb,  206.9) 

The  first  two  elements  of  this  family  are  entirely  non-metallic,  while 
the  others  are  metals  exhibiting  properties,  however,  which  reveal 
resemblances  to  the  non-metals.  The  group,  for  advantage  of  study, 
may  be  and  is  frequently  divided  into  a  primary  group  including  carbon 
and  silicon,  and  a  secondary  group  composed  of  germanium,  tin  and  lead. 
All  of  the  elements  of  the  family  possess  a  maximum  valence  of  four. 
They  unite  with  four  atoms  of  hydrogen  or  of  the  halogens.  With  the 
exception  of  silicon,  they  also  form  compounds  in  which  they  are  bivalent. 

PRIMARY    GROUP. 

Carbon  and  silicon  resemble  one  another  in  their  chemical  conduct. 
Some  of  their  physical  properties  are  quite  similar.  Their  derivatives 
are  also  very  much  alike — the  halogen  compounds  exhibiting  similar 
properties,  the  oxides  being  weak  acid  anhydrides  which  are  capable 
of.  forming  stable  salts  with  bases.  The  acids  of  carbon  and  silicon  are 
unstable,  decomposing  into  water  and  the  acid  anhydride.  The  normal 
valence  of  the  two  elements  is  four. 

The  essential  difference  in  the  chemical  conduct  of  carbon  and  silicon 
is  the  ability  of  atoms  of  the  former  to  combine  with  each  other  and  form 
"chain  compounds,"  i.e.,  carbon  possesses  the  property  of  satisfying  its 
own  bonds  of  valence.  This  property  is  indicated  in  the  following 
structural  formulae  of  ethyl  alcohol  (C2H5OH),  propane  (C3H8),  and 
butane  (C4H10): 

H      H  H        H      H 

II  III 

H— C  —  C  —  O  —  H,       H— C  —  C  —  C  —  H, 

I         I  I 

H      H  H       H      H 

H       H       H       H 

I         I         I         I 
H—  C  —  C  —  C  —  C—  H. 

I         I  I 

H       H       H       H 

226 


CARBON    FAMILY.  227 

Silicon  does  not  possess  the  power  of  satisfying  its  own  bonds,  con- 
sequently it  does  not  form  "chain  compounds." 

CARBON,    C. 

At.  Wt.  12.00     Mol.  Wt.  (Cx). 

Carbon  occurs  in  the  free  and  almost  pure  state  in  nature  in  several 
allotropic  forms  known  as  diamond,  graphite  or  plumbago  and  amorphous 
carbon.  It  is  also  found  in  the  combined  condition  in  all  living  things, 
in  carbon  dioxide  and  carbonates,  in  turf,  peat,  lignite  or  brown  coal, 
bituminous  coal  and  anthracite  coal,  and  in  such  mineral  oils  as  asphal- 
tum  and  petroleum.  As  the  result  of  the  power  of  carbon  to  satisfy 
its  own  bonds,  the  number  of  possible  carbon  compounds  is  so  large 
that  it  appears  necessary,  for  purposes  of  convenience,  to  treat  them 
apart  in  a  separate  portion  or  special  phase  of  chemistry.  With  hydrogen 
carbon  forms  practically  an  unlimited  number  of  compounds  into  which 
nearly  all  other  elements,  especially  oxygen,  nitrogen,  the  halogens  and 
sulphur  can  enter.  The  derivatives  of  carbon  have  been  termed  organic 
compounds  because  of  the  idea  which  prevailed  for  many  years,  that 
these  compounds  could  not  be  produced  without  the  intervention  of 
life,  i.e.,  that  their  artificial  production  was  impossible.  On  account 
of  this  original  belief  the  chemistry  of  the  carbon  compounds  is  commonly 
known  as  organic  chemistry. 

Although  it  is  true  tnat  a  large  number  of  the  derivatives  of  carbon 
are  obtained  exclusively  from  animal  and  vegetable  organisms,  yet  many 
of  them  are  prepared  artificially  from  simpler  ones  or  from  the  elements 
by  simple  synthetic  methods.  The  student  is  also  reminded  of  the  fact 
that  hundreds  of  definite  chemical  compounds,  including  drugs  and 
dyes  of  great  value,  apparently  unknown  to  either  animal  or  vegetable 
life,  have  been  prepared  artificially.  The  preparation  of  many  other 
compounds  is  delayed  solely  because  of  their  unusual  complexity  and 
instability. 

Experiment  I. — Properties  of  Carbon. 

(a)  Enumerate  the  known  allotropic  modifications  of  this  element. 
Examine   specimens   of   each.     Which   represents   the   purest   carbon? 
Tabulate  their  principal  physical  properties  ? 

(b)  Preparation  of  amorphous  carbon. 

(bi)  Place  about  i  gram  of  cane  sugar  (C12H22On)  in  an  old  test  tube 
and  heat  until  vapors  cease  to  appear.  Was  water  liberated  during 
the  process?  Pulverize  the  black  residue  in  a  mortar.  What  is  it? 
Equation  ? 

(62)  Cover  the  bottom  of  a  crucible,  preferably  an  iron  crucible 
provided  with  lid  and  delivery  tube,  with  sand.  Place  a  number  of  small 
pieces  of  soft  pine  (largely  cellulose,  (C6H10O5)n  in  the  crucible  and 
cover  with  sand  to  partially  exclude  the  air.  Support  the  crucible  in  a 


228  EXPERIMENTAL  CHEMISTRY. 

pipe-stem  or  wire  triangle  and  heat  for  about  a  half-hour  or  until  the 
contents  of  the  crucible  cea»e  to  smoke.  Allow  the  crucible  to  cool, 
then  pour  the  contents  out  upon  an  iron  plate.  Place  one  or  two  of  the 
charred  pieces  of  wood  in  a  mortar.  It  should  reduce  readily  to  a  fine 
powder.  What  is  it?  Equation? 

(63)  Close  the  holes  at  the  bottom  of  a  lighted  Bunsen  burner  and 
hold  a  piece  of  cold  glass  tubing  in  the  upper  portion  of  the  flame.     What 
is  the  black  deposit  ?     By  what  two  names  is  it  known  ? 

Repeat,  using  the  flame  of  a  candle  ( ?). 

(64)  (Hood).     Apply  a  flame  to  a  piece  of  camphor  gum  (C10H1CO) 
about  the  size  of  a  pea.     Does  it  burn  with  a  "sooty"  flame?     Explain. 

(65)  Name   the   different   kinds   of   charcoal.     Give   an   example   of 
"destructive    distillation,"    and    define    same.     What    is    boneblack? 
Coke? 

(c)  Porosity  of  charcoal.     Observe  that  a  piece  of  ordinary  charcoal 
floats  upon  the  water.     Place  a  small  piece  of  it  in  a  test  tube  half  full  of 
water;  by  means  of  a  long  glass  rod  or  a  piece  of  wire  force  the  charcoal 
down  to  the  bottom  of  the  tube  and  boil  the  water  for  several  minutes 
or  until  the  charcoal  shows  little  or  no  tendency  to  float.     Explain. 

(d)  Decolorizing  action  of  charcoal.     To  a  test  tube  one-fourth  full 
of  powdered  animal  charcoal  add   10  cm.3   of  water  which  has  been 
colored  by  the  addition  of  a  few  drops  of  an  indigo  solution  or  a  litmus 
solution;  shake  thoroughly  for  a  few  minutes,  then  heat  to  boiling;  filter 
through  a  wet  filter  paper  and  compare  the  filtrate  with  that  of  the  orig- 
inal solution  (?).     Explain. 

Note. — A  solution  of  brown  sugar  may  be  substituted  for  either  of 
the  above  solutions. 

(e)  Deodorizing  action  of  charcoal.     Fill  a  test  tube  nearly  one-third 
full  of  powdered  charcoal;  add  5  cm.3  of  a  weak  solution  of  hydrogen 
sulphide;  cork  the  tube  tightly  and  shake  vigorously;  set  the  tube  aside 
and  after  fifteen  to  thirty  minutes  note  the  odor  of  the  contents  (?). 
Repeat  until  your  results  are  definite.     Explain. 

(/)  Reducing  action  of  charcoal.  Mix  thoroughly  on  a  piece  of  paper 
2  grams  of  copper  oxide,  CuO,  and  .5  grams  of  powdered  charcoal.  Place 
the  mixture  in  a  hard  glass  test  tube  (ignition  tube) ;  clamp  the  tube  in  a 
nearly  horizontal  position  and  heat  strongly.  The  test  tube  may  be 
provided  with  a  cork  and  delivery  tube  which  should  dip  into  a  solution 
of  calcium  hydroxide,  Ca(OH)2,  in  a  test  tube,  or  this  second  test  tube 
containing  the  Ca(OH)2  solution  may  be  held  vertically  under  the  mouth 
of  the  other  tube  so  that  the  heavy  gas  which  is  evolved  when  the  mixture 
is  heated  may  fall  into  the  tube,  then  close  the  tube  with  the  thumb  and 
shake.  Explain  the  milky  appearance  of  the  solution.  Examine  the 
residue  in  the  ignition  tube.  This  may  be  done  by  rubbing  it  in  the 
mortar  and  washing  away  the  lighter  particles.  Results?  Explain. 
Equation  ? 

Recall  the  interaction  of  As-jOg  and  charcoal  when  heated  together  ( ?). 


CARBON    FAMILY. 


229 


Equation?  What  is  such  a  type  of  reaction  called?  Is  carbon  a 
chemically  active  substance  at  ordinary  temperatures  ?  How  does  the 
charring  of  wood  preserve  it? 

Carbon  and  Oxygen. 

Experiment  II. — Preparation  and  Properties  of  Carbon  Dioxide. 

(a)  Recall  or  repeat  Exp.  III.  "Oxygen"  (?).  Do  carbon  and  oxygen 
react  at  ordinary  temperatures?     What  is  formed  by  their  interaction? 
Equation  ?     What  is  the  usual  test  for  carbon  dioxide  ? 

(b)  Show  that  CO2  k  formed  when  ordinary  combustibles,  such  as 
wood,  paper,  illuminating  gas,  etc.,  are  burned  in  air. 


FIG.  40. — Kipp  Generator. 


(c)  Hold  a  piece  of  charcoal  in  the  Bunsen  flame  and  describe  its 
combustion.     Products?     Repeat     experiment     using    graphite     (pencil 
lead)  (?).     Define  combustion. 

(d)  Pour  a  few  cm.3  of  clear  limewater  (Ca(OH2))  into  a  test  tube 
and  by  means  of  a  glass  tube  blow  air  from  the  lungs  through  the  liquid. 
Result  ?     Explain.     Would  you  infer  that  carbon  undergoes  combustion 
within  the  human  body? 

(e)  Usual  laboratory  method  of  preparation  (Fig.  40).    Put  several  small 
pieces  of  marble  (calcium  carbonate,  CaCO3)  in  the  hydrogen  generator 
and  add  enough  water  to  seal  the  bottom  of  the  thistle  tube.     Connect 
the  generator  with  a  wash  bottle  half  filled  with  strong  sulphuric  acid. 
Pour  strong  hydrochloric  acid  into  the  flask,  as  needed,  through  the 


230  EXQERIMENTAL  CHEMISTRY. 

thistle  tube.  Collect  five  or  six  bottles  of  the  gas  by  displacement  of 
air,  placing  the  bottle  with  the  mouth  upward.  Prove  that  the  evolved 
gas  is  carbon  dioxide.  Note  its  color,  odor  and  taste.  Is  the  gas  com- 
bustible? Does  it  support  combustion?  Pour  one  bottleful  of  the  gas 
very  slowly  into  another  bottle  of  about  the  same  size  containing  air. 
Lower  a  lighted  splinter  into  the  second  bottle.  Result  ?  Conclusion  ? 
Is  the  gas  heavier  or  lighter  than  air  ? 

Counterpoise  a  beaker  on  the  balance,  then  pour  carbon  dioxide 
into  it  (?). 

Filter  the  contents  of  the  generating  flask;  evaporate  (hood)  a  portion 
of  the  filtrate  to  dryness;  heat  the  residue  until  fumes  of  hydrochloric 
acid  cease  to  be  evolved.  Dissolve  a  portion  of  the  residue  in  distilled 
water  and  test  for  a  chloride.  Test  for  calcium  by  dipping  a  clean 
moist  platinum  wire  into  the  solid  residue  and  holding  it  in  the  Bunsen 
flame.  A  yellowish-red  flame  indicates  the  presence  of  calcium.  What 
is  the  name  of  the  salt  (residue)  in  the  evaporating  dish?  Indicate  by 
equations  the  reactions  involved  in  the  preparation  of  carbon  dioxide. 

(/)  Try  the  action  of  hydrochloric  acid  upon  small  portions  of  each 
of  the  following  salts:  sodium  carbonate  (Na^COJ,  sodium  hydrogen 
carbonate  (NaHCO3),  potassium  carbonate  (K2CO3),  barium  carbonate 
(BaCO3).  Results?  Is  CO2  evolved  in  each  case?  (Use  the  "loop 
tube "  to  carry  a  film  of  lime-water  to  the  mouth  of  the  test  tube.  This 
renders  the  testing  for  carbon  dioxide  a  very  simple  process.) 

(g)  Place  about  a  gram  of  sodium  carbonate  or  powdered  magnesite 
(MgCO3)  in  a  test  tube,  and  heat  strongly.  Is  carbon  dioxide  evolved? 

Interpret  the  following  equations: 

C  (charcoal)  +  O2  — >  CO2  +  96,980  cal. 

C  (retort  carbon)  +  O2  — >  CO2  +  96,530  cal. 

C  (graphite)       +  O2  — »  CO2  +  93,360  cal. 

C  (diamond)      +  O2  — >  CO2  +  93,240  cal. 

Experiment  III. — Optional.     Synthesis  of  Carbon  Dioxide. 

This  operation  may  be  carried  out  by  drawing  simultaneously  purified 
air  through  a  tube  over  hot  charcoal,  and  the  product  into  a  solution  of 
calcium  or  barium  hydroxide..  (Instructions.) 

Experiment  IV. — Carbonic  Acid.     Carbonates. 

(a)  Half  fill  a  clean  bottle  with  distilled  water;  test  the  water  with 
litmus  paper  or  a  few  drops  of  a  solution  of  phenolphthalein  ( ?).  Allow 
a  stream  of  carbon  dioxide  to  bubble  through  the  water  until  the  latter 
is  saturated.  Test  the  water  with  litmus  paper.  Results?  Con- 
clusions? Ascertain  the  relative  conductivity  of  the  solution  (?).  CO2 
is  the  anhydride  of  what  acid?  (This  acid,  carbonic  (H2CO3),  has 
never  been  prepared  in  the  free  state.)  Is  the  acid  stable?  If  not, 
into  what  substances  does  it  decompose  ?  What  ions  does  the  acid  yield  ? 
Is  it  a  "  strong"  acid  ?  State  the  reasons  for  your  answer. 


CARBON    FAMILY.  231 

Can  you  explain  why  carbonic  acid  is  so  readily  liberated  from  car- 
bonates? State  Berthollet's  law.  Henry's  law. 

(CO2l  Aq.)  =  5800  cal. 

(b)  Conduct  carbon  dioxide  into  lime-water,  or  baryta  water  (Ba(OH)2) 
until  the  gas  ceases  to  be  absorbed.  Results?  Filter  and  test  the 
precipitate  for  carbon  dioxide  (?).  What  salt  is  formed  by  the  action 
of  CO2  on  lime-water?  Is  it  soluble  in  water?  In  acids?  Equations? 

Saturate  a  solution  of  sodium  hydroxide  with  carbon  dioxide.  Set 
the  solution  aside  and  allow  the  liquid  to  evaporate  spontaneously. 
Prove  that  the  residue  is  a  carbonate.  Equations  ? 

What  class  of  salts  are  yielded  by  carbonic  acid  (H2O  +  CO2)  ?  What 
is  the  basicity  of  carbonic  acid?  How  would  you  determine  whether  a 
rock  specimen  contained  a  carbonate  or  not  ? 

Experiment  V. — Formation  of  Calcium  Acid  Carbonate. 

Conduct  carbon  dioxide  into  lime-water  until  the  precipitate  which 
forms  at  first,  disappears.  Filter.  Divide  the  filtrate  into  two  portions. 
To  one  portion  add  clear  lime-water  (?).  Heat  the  second  portion  in 
a  test  tube  (  ?).  Why  did  the  precipitate  disappear  in  the  first  reaction  ? 
Why  does  it  appear  after  heating  ?  Equations  ?  Explain  the  formation 
of  the  incrustations  on  the  inside  of  tea-kettles  and  steam  boilers. 

All  normal  carbonates  except  those  of  the  alkalies  are  insoluble  in 
water.  They  are  decomposed  by  hydrochloric  acid  with  an  evolution 
of  carbon  dioxide. 

Experiment  VI. — Optional.  (Quant.)  Estimation  of  Carbon  Dioxide  in 
a  Carbonate. 

(a)  Report  a  method  for  the  determination  of  carbon  dioxide  in  a 
carbonate  in  which  the  evolved  gas  is  absorbed  in  a  weighed  apparatus. 

(b)  Estimation    by   difference.     The   epitome    of   the    process   is    as 
follows:     A  known  weight  of  the  carbonate  is  decomposed  by  dilute 
acid  in  a  weighed  apparatus  and  the  carbon  dioxide  is  dried  as  it  escapes 
or  is  expelled  through  a  bulb  containing  strong  sulphuric  acid.     The 
loss  in  weight  which  results  represents  the  carbon  dioxide  in  the  sub- 
stance being  examined. 

Schrotter's  apparatus  is  a  convenient  form  of  the  apparatus  required 
for  the  performance  of  this  experiment.  Draw  a  diagram  of  the  appa- 
ratus and  explain  its  manipulation.  (See  model  in  laboratory  or  text 
books  on  analytical  chemistry.  Figures  of  various  forms  of  the  apparatus 
usually  appear  in  catalogues  of  laboratory  supplies.) 

Experiment  VII. — (Quant.)  Determination  of  the  Density  of  Carbon 
Dioxide.  Molecular  Weight. 

Provide  a  300  cm. 3  flask  with  a  tightly  fitting  rubber  cork.  See  that 
the  apparatus  is  perfectly  clean  and  dry.  Weigh  the  flask  and  cork. 


232  EXPERIMENTAL  CHEMISTRY. 

Record  this  weight  as  the  weight  of  flask,  cork,  and  air  which  filled  flask. 
Fill  the  flask  with  carbon  dioxide  by  displacement  of  air;  stopper  the 
flask  and  weigh.  Record  weight.  Repeat  the  operation  of  filling  the 
flask  with  carbon  dioxide  and  weighing,  until  the  weight  becomes  practi- 
cally constant.  Calibrate  the  flask  as  in  previous  experiments.  The 
weight  of  the  empty  flask  is  found  by  subtracting  from  the  weight  of  the 
cork  and  flask  filled  with  air  the  weight  of  a  volume  of  air  equal  to  the 
capacity  of  the  flask.  (Air  under  normal  conditions,  i  cm.3  =  .00129 
gram.)  What  is  the  weight  of  the  carbon  dioxide  in  the  flask?  Of  i 
litre  of  the  gas? 

Calculate  the  molecular  weight  of  carbon  dioxide  from  the  experi- 
mental data  (?).  What  is  your  percentage  of  error? 

Mention  sources  of  error  in  determining  the  density  of  the  gas  by  above 
method. 

Experiment  VIII. — Preparation  and  Properties  of  Carbon  Monoxide. 

(a)  Assemble  apparatus  similar  to  that  which  you  used  for  the  pre- 
paration of  chlorine.     Put  10  grams  of  oxalic  acid  (C2H2O4)  in  the  gener- 
ating flask.     Half  fill  the  wash  bottle  with  lime-water  or  a  caustic-soda 
solution.     Pour  35   cm.3   of  concentrated   sulphuric   acid   through  the 
thistle  tube  into  the  flask;  heat  the  latter  gently.     Collect  three  or  four 
bottles  of  the  evolved  gas  over  water.     Add  a  few  drops  of  lime-water 
to  a  bottle  of  the  gas;  place  your  hand  over  the  mouth  of  the  bottle  and 
shake  vigorously.     Is  there  any  evidence  that  carbon  dioxide  is  present  ? 
Set  fire  to  a  bottle  of  the  gas  and  notice  the  characteristic  blue  flame. 
After  the  gas  has  burned,  pour  a  few  cm.3  of  lime-water  into  the  bottle; 
shake.     Results  ?     Describe  any  visible  changes  that  may  have  occurred 
in  the  wash-bottle.     Add  a  few  drops  of    acid  to  the  solution  and  test 
for  carbon  dioxide  (?).     What  gas  besides  carbon  monoxide  was  evi- 
dently evolved  by  the  action  of  H2SO4  on  C2H2O4?         This  chemical 
action  is   due   largely   to  what  particular  property  of    sulphuric  acid? 
Equations?  May  CO  act  as  a  reducing  agent?     Explain.     Interpret  the 
following  equations  (?). 

(b)  Optional.     Intimately  mix  2  grams  of  zinc  oxide   (ZnO)   and   i 
gram  of  powdered  charcoal.     Introduce  the  mixture  into  a  hard  glass 
test  tube  and  heat  strongly.     Apply  a  flame  to  the  mouth  of  the  tube. 
Result  ?     Equations  ? 

(C,  O)  =  28,500  cal.         (CO,O)  =  68,400  cal. 

Cu,O  (37,100  cal)  +  C,O  (28,500  cal.)— >  Cu  +  C,O2  (96,900  cal.) 
+  31,300  cal. 

"  If  an  element  combines  with  another  according  to  multiple  pro- 
portions, there  usually  occurs,  in  the  union  of  the  first  atom,  a  greater 
disengagement  of  heat  than  with  the  following  atom  (compare  nitrogen 
oxides).  The  numbers  above,  on  the  contrary,  show  that  the  union 
of  the  second  atom  of  oxygen  wuth  carbon  (CO,O)  sets  free  68,400 
calories;  that  of  the  first  atom  (C,O),  however,  only  28,500  calories. 


CARBON   FAMILY.  233 

This  can  only  be  explained  by  the  fact  that,  for  the  vaporization  and 
disaggregation  of  the  solid  carbon  molecules,  heat  is  necessary.  If  we 
assume  that  the  direct  union  of  the  first  atom  also  disengaged  68,400 
calories,  it  would  follow  from  this  that,  in  the  dislocation  of  twelve 
parts  carbon  (amorphous)  by  weight  into  gaseous  free  atoms,  39,900 
(68,400-28,500)  calories  were  absorbed."  Richter. 

The  above  equations,  indicating  the  heat  disengagement  for  each 
stage  in  the  oxidation  of  carbon,  furnish  an  excellent  illustration  of  the 
Law  of  Hess.  (State  the  law.) 

Experiment  IX. — (Quant.)  Determination  of  the  Density  of  Carbon 
Monoxide.  Molecular  Weight. 

The  density  of  carbon  monoxide  may  be  determined  by  a  method 
identical  with  that  recorded  in  Exp.  III.  " Nitrogen  and  the  Atmosphere.'' 

Carbon  and  Hydrogen  (Hydrocarbons). 

Experiment  X. — Preparation  and  Properties  of  Methane  (Marsh  Gas). 

Heat  a  few  grams  of  sodium  acetate  in  a  porcelain  dish  until  the  water 
of  crystallization  has  been  expelled.  Heat  a  mixture  consisting  of 
equal  parts  of  dehydrated  sodium  acetate  and  soda  lime  in  a  hard  glass 
test  tube  clamped  in  a  nearly  horizontal  position.  The  tube  should  be 
fitted  with  a  cork  and  delivery  tube.  Collect  the  gas  over  water.  Note  its 
odor,  color  and  taste.  Does  it  burn  ?  Does  the  gas  explode  when  mixed 
with  air  ?  Attach  a  jet  to  the  delivery  tube  and  burn  the  gas  when  all  the 
air  has  been  displaced  from  the  apparatus.  Note  the  degree  of  luminosity 
of  the  flame.  Is  water  formed  during  the  burning  of  the  gas?  Test. 
Carbon  dioxide  ?  Test.  Equations  ?  What  is  the  name  and  molecular 
formula  of  the  gas? 

(C,*  H4)  =  21,700  cal.         (C2,  H6)  =  28,600  cal. 

Experiment  XI. — (Optional.)  Preparation  of  Ethylene  by  Elimination 
of  Water  from  Alcohol.  Ethylene  Bromide. 

The  following  method  is  suggested  by  Gattermann:  A  mixture  of 
25  grams  of  alcohol  and  150  grams  of  concentrated  H2SO4  is  heated, 
not  too  strongly,  in  a  liter  round  flask  on  a  wire  gauze  covered  with  thin 
asbestos,  or  sand  bath.  As  soon  as  an  active  evolution  of  ethylene 
takes  place,  add,  through  a  dropping  funnel,  a  mixture  of  one  part 
alcohol  and  two  parts  concentrated  H2SO4  (made  by  pouring  300  grams 
of  alcohol  into  600  grams  of  sulphuric  acid,  with  constant  stirring), 
slowly,  so  that  a  regular  stream  of  gas  is  evolved.  If  the  mixture  in  the 
flask  forms  badly  with  a  separation  of  carbon,  it  has  been  too  strongly 
heated,  and  it  is  advisable  to  empty  the  flask  and  begin  the  operation 
anew.  In  order  to  free  ethylene  from  alcohol,  ether  and  sulphur  dioxide, 
it  is  passed  through  a  wash  bottle  containing  sulphuric  acid,  and  a  second 

#  Amorphous  carbon. 


234 


EXPERIMENTAL  CHEMISTRY. 


one,  provided  with  three  tubulures,  the  central  one  supplied  with  a  safety- 
tube,  containing  a  dilute  solution  of  caustic  soda.  It  is  well  to  use  a 
third  wash  bottle  containing  sulphuric  acid.  When  the  air  in  the 
apparatus  has  been  displaced,  fill  a  small  narrow-mouthed  bottle  provided 
with  a  greased  glass  stopper,  with  the  gas  by  downward  displacement 
of  air.  Allow  a  drop  of  bromine  to  fall  into  the  bottle,  then  replace  the 
stopper  quickly.  Observe  any  changes  taking  place  within  the  bottle 
After  a  few  minutes  remove  the  stopper  under  water.  Results  ?  Explain. 
Attach  a  jet  to  the  delivery  tube  and  burn  the  gas.  Compare  the 
degree  of  luminosity  of  the  flame  with  that  of  methane  (?).  Equations? 

(C2,  H4)  =        -  2700  cal. 
Experiment  XII. — Preparation  and  Properties  of  Acetylene. 

Fit  a  test  tube  with  a  single-hole  stopper  carrying  a  piece  of  straight 
glass  tubing  about  5  cm.  long  and  ending  flush  with  the  lower  side  of  the 
stopper.  Fill  the  test  tube  nearly  full  of  water  then  drop  into  it  a  small 
piece  of  calcium  carbide,  CaC2;  insert  the  stopper.  Note  the  odor  of 


FIG.  41. — The  Manufacture  of  Coal  Gas. 

the  evolved  gas.  When  the  air  in  the  tube  has  been  displaced,  light 
the  gas  which  is  issuing  from  the  jet.  Describe  the  character  of  the 
flame.  Hold  a  glass  plate  just  above  the  flame  (?).  Would  you  infer 
that  there  is  perfect  combustion  of  carbon  ?  Why  ?  What  is  the  name 
and  molecular  formula  of  the  gas?  Equations  indicating  the  action  of 
the  carbide  on  water  and  the  combustion  of  the  gas?  Will  air  burn  in 
an  atmosphere  of  the  gas?'  Can  you  offer  a  possible  explanation  as  to 


CARBON    FAMILY.  235 

why  C2H2  and  C2H4  are  endothermic  compounds,  and  CH4  and  C2H6 

are  exothermic?     (See  carbon  monoxide.)     (C2H2)  =     — 47,800  cal. 

Coal  Gas  is  formed  by  the  destructive  distillation  (Fig.  41)  of  bituminous 
coal.  The  composition  of  the  products  obtained  depend  largely  upon  the 
nature  of  the  coal  and  the  process  of  distillation.  These  products  are:  (i) 
coke;  (2)  coal  tar;  (3)  gas  liquor  containing  ammonia  and  other  products, 
and  known  as  ammoniacal  liquor;  and  (4)  coal  gas.  The  coal  gas  is 
a  mixture  of  various  gases.  The  components  of  the  mixture  may  be 
divided  into  three  classes  as  follows: 

Illuminants.  Diluents.  Impurities. 

(About  6%)  (About  90%)  (About  3%) 

Ethylene,     C2H4  Hydrogen,         H2          Nitrogen,        N2 

Propylene,  C3He  Methane,        CH4          Carbon  dioxide  CO2 

Butylene      C4H8  Carbon  monoxide  CO  Hydrogen  sulphide,  H2S 
Acetylene    C2H2  Ammonia,     NH3 

Allylene,  C3H4 
Benzene,  C6H0 
Heavy  hydrocarbons. 

Some  of  the  impurities  are  removed  entirely  from  the  purified  gas; 
i.e.,  before  the  gas  is  delivered  to  the  gas  mains.  The  gas  is  often  called 
"illuminating  gas."  The  student  is  referred  to  Sadtler's  Hand-book 
of  Industrial  Organic  Chemistry  and  Thorpe's  Outlines  of  Industrial 
Chemistry. 

Experiment  XIII. — Preparation  and  Properties  of  Illuminating  Gas. 

Fill  a  hard  glass  test  tube  one-half  full  of  coarsely  powdered  bituminous 
or  cannel  coal;  place  a  plug  of  glass  wool  or  shredded  asbestos  above 
the  coal  to  hold  it  in  place;  clamp  the  tube  in  a  horizontal  position  and 
connect  it  with  an  empty  wash  bottle  which  communicates  with  a  U-tube 
containing  strips  of  litmus  paper  in  one  limb  and  filter  paper  moist- 
ened with  lead  acetate  in  the  other  limb;  a  third  wash  bottle  is  one-third 
full  of  lime-water.  Heat  the  ignition  tube  gently  with  the  Bunsen  flame. 
As  soon  as  all  of  the  air  has' been  driven  out  of  the  apparatus  collect  two 
bottles  of  the  gas  by  water  displacement.  Disconnect  the  apparatus. 
Examine  the  gas  in  the  bottles,  noting  its  color,  odor,  inflammability 
and  color  of  flame.  Of  what  does  the  gas  consist  ?  Does  the  residue 
in  the  ignition  tube  suggest  some  form  of  carbon  ?  What  name  is  applied 
to  it  ?  Describe  in  full  and  explain  any  changes  which  may  have  occurred 
within  the  wash-bottles.  What  are  the  products  of  the  combustion  of 
the  chief  components  of  illuminating  gas?  Devise  an  experiment  to 
prove  your  conclusions.  (Hint. — Use  gas  from  jet.) 

Make  a  brief  statement  of  the  composition  and  manufacture  of 
"producer  gas."     Of  "water  gas." 

Carbon  and  Sulphur. 


236  EXPERIMENTAL  CHEMISTRY. 

Experiment  XIV.  —  Properties  of  Carbon  Bisulphide. 

Procure  a  few  cm.3  of  carbon  disulphide  in  a  test  tube  from  the  side- 
shelf  reagent.  Note  its  odor,  color  and  high  refractive  index.  What 
is  its  specific  gravity  ?  Place  a  drop  of  the  liquid  on  the  hand  and  force 
air  over  it  (  ?).  Place  a  very  small  quantity  of  the  substance  in  an  evapo- 
rating dish,  then  bring  the  heated  end  of  a  glass  rod  near  the  surface  of 
the  liquid  (?).  Would  you  infer  that  carbon  disulphide  is  very  inflam- 
mable ?  Is  it  miscible  with  water  ?  With  alcohol  ?  It  has  been  used  in 
previous  experiments  as  a  solvent  for  what  substances?  Is  rubber 
soluble  in  it  ?  If  so,  allow  the  liquid  to  evaporate  spontaneously.  Result  ? 
Enumerate  other  uses  of  carbon  disulphide.  What  is  the  molecular 
formula  of  carbon  disulphide  ?  Is  it  an  endothermic  compound  ?  (See 
following  equation.)  What  would  you  infer  as  to  its  stability? 

(C,  S2  —  liq.)  =  —  19,600  cal. 

Carbon  and  the  Halogens.  —  Although  carbon  does  not  combine  di- 
rectly with  the  members  of  the  halogen  family,  halides  of  carbon  are 
formed  by  the  action  of  the  halogens  on  the  hydrocarbons.  (Recall  Exp. 
XL)  In  our  previous  work  we  have  observed  that  chlorine  is  capable 
of  withdrawing  or  substituting  itself  for  the  hydrogen  of  water,  hydrogen 
sulphide,  ammonia,  etc.  The  halogens  react  very  similarly  with  the 
hydrocarbons  forming,  by  a  process  termed  substitution,  a  series  of 
carbon  compounds  known  as  substitution  products.  By  such  reactions 
the  following  familiar  substances  are  obtained:  Methyl-chloride  (CH3C1), 
chloroform  (CHC13),  iodoform  (CHI3),  carbon  tetrachloride  (CC14), 
etc.  The  latter  (CC14)  is  a  colorless  liquid  having  a  peculiar  odor. 
It  has  a  sp.  gr.  of  1.6  and  boils  at  77°.  It  is  the  final  product  of  the 
action  of  chlorine  upon  methane  (CH4)  or  upon  chloroform. 

CH4  +    C12  —  CH3C1   +  HC1 
CH4  +  2C12—  CH2C12  +  2HC1 


CH4  +  4C12  —  CC14      +  4HC1 

(C,  C14  —  gas)  =  21,000  cal.         (C,  C14—  liq.)  =  28,300  cal. 
(C2,  C14  —  gas)  =  —  1  100  cal. 

The  heats  of  formation  of  the  chlorides  of  carbon  approximate  those 
of  the  hydrogen  derivatives,  indicating  that  the  affinity  of  the  two  ele- 
ments for  carbon  is  practically  the  same. 

Carbon  and  Nitrogen.  —  These  two  elementary  substances  apparently 
do  not  possess  any  direct  affinity  for  one  another,  therefore  the  union 
of  the  two  is  usually  effected  by  indirect  processes.  However,  small 
quantities  of  cyanogen  (CN  or  Cy)  are  formed  when  electrical  dis- 
charges take  place  between  carbon  poles  in  an  atmosphere  of  nitrogen. 
The  production  is  facilitated  if  one,  at  least,  of  the  elements  is  in  the 
form  of  a  compound.  Cyanogen  is  an  endothermic  compound  —  a 
colorless  gas  with  an  odor  resembling  that  of  peach  blossoms,  or  better, 


CARBON    FAMILY.  237 

almonds.  It  burns  with  a  characteristic  violet-blue  flame,  or  more 
accurately,  with  a  "pink  flame  edged  with  green."  Its  sp.  gr.  is  1.805 
(Air  =  i).  Nascent  cyanogen  is  recognized  as  a  radical  or  group  with 
the  formula,  CN.  It  is  seen  at  once  from  its  structure  ( — C  EEN)  that 
it  is  an  unsaturated  compound — a  negative  univalent  radical.  Like 
other  univalent  groups  (OH,  NH2,  CH3),  it  cannot  exist  in  the  free  state, 
but  combines  with  itself  to  form  a  double  molecule  of  free  cyanogen  (C2N2), 
known  as  dicyanogen.  The  structural  formula  (N  ==  C  —  C  EE  N) 
shows  carbon  satisfying  its  own  bonds. 

In  its  chemical  relations  carbon  exhibits  the  mutability  peculiar  to 
organic  groups.  It  combines  with  hydrogen  under  the,,  influence  of  a 
quiet  electric  discharge  to  form  the  very  "weak"  but  extremely  poisonous 
hydrocyanic  or  Prussic  acid  (HCN).  This  acid  is  a  colorless  mobile 
liquid  which  boils  at  about  27°. 

An  aqueous  solution  possesses  small  electrical  conductivity.  It  forms 
cyanides  by  exchanging  its  hydrogen  for  metals.  These  salts,  with  the 
exception  of  potassium  cyanide  which  undergoes  hydrolysis  to  a  marked 
degree,  are  relatively  stable  and  uniformly  poisonous.  The  extreme 
feebleness  of  the  acid  is  shown  by  the  fact  that  the  carbon  dioxide  of  the 
air  liberates  it.  For  this  reason  the  salt  usually  possesses  the  odor  of 
hydrocyanic  acid. 

The  tendency  of  the  salts,  like  potassium  cyanide  (KCN),  in  the 
presence  of  easily  reducible  and  highly  heated  oxides  to  form  cyanates 
(KCNO)  indicates  that  hydrocyanic  acid  is  an  unsaturated  compound, 
and  possesses  a  structure  similar  to  that  indicated  by  the  following 
formula,  H  —  N  =  C  =  .  This  behavior  of  potassium  cyanide  accounts 
for  its  frequent  use  as  a  suitable  flux  for  securing  reductions.  It  mani- 
fests a  similar  tendency  to  combine  with  sulphur,  forming  a  salt  belonging 
to  a  class  known  as  the  thiocyanates,  e.g.,  KCNS. 

Cyanogen  also  unites  with  iron  to  form  a  compound  or  rather  a  group 
which  combines. with  various  other  metals  to  produce  two  series  of  salts, 
namely,  the  jerrocyanides  and  the  jerricyanides.  The  potassium  salt 
of  the  former,  potassium  ferrocyanide  (K4Fe(CN)6.3Aq.),  has  been 
given  the  name  of  yellow  prussiate  oj  potash;  the  corresponding  salt  of 
the  latter  is  known  as  potassium  ferricyanide,  K3Fe  (CN)6  or  K6FeCy6. 

(C,  N,  H— gas)  =     -27,500  cal. 
Experiment  XV. — Preparation  and  Properties  of  Cyanogen. 

Note. — This  gas  is  intensely  poisonous  and  should  not  be  prepared 
by  the  student. 

Cyanogen  may  be  prepared  by  (i)  heating  the  cyanides  of  mercury 
or  silver;  (2)  the  dry  distillation  of  ammonium  oxalate,  and  (3)  the  fusion 
of  potassium  cyanide  and  mercuric  chloride. 

Hg(CN)2—  Hg  +  "QNL 
(NH4)2C204  ->  4H20  +  C2N2 
2KCN  +  HgCl2  —  Hg  +  2KC1  +  C2N2 


238  EXPERIMENTAL  CHEMISTRY. 

Experiment  XVI. — Preparation  and  Properties  of  Potassium  Cyanide. 
Potassium  Thiocyanate. 

(a)  Potassium  cyanide  (KCN)  may  be  prepared  by  heating  a  nitrogen- 
ous carbon  compound  like  gelatin  (C^H^N^O^,  approx.),  with  a 
small  piece  of  potassium.  The  KCN  which  forms  is  extracted  with 
water. 

To  a  few  cm.3  of  a  silver  nitrate  solution  add  a  few  drops  of  a  solution 
of  KCN.  Results  ?  Add  an  excess  of  KCN.  Result  ?  Equations  ? 

(6)  A  test  for  cyanogen  compounds.  To  a  dilute  aqueous  solution  of 
KCN  add  equal  quantities  of  a  ferrous  and  a  ferric  salt.  Result  ?  Add 
an  excess  of  a  potassium  hydroxide  solution,  then  acidify  with  HC1. 
Results?  Before  attempting  to  state  the  composition  of  the  final  pro- 
duct, add  a  few  drops  of  potassium  ferrocyanide  (K4FeC6N6)  to  a  solution 
of  a  ferric  salt.  Result?  What  are  your  conclusions  now  as  to  the 
composition  of  the  product  referred  to  above  ? 

(c)  Potassium  thiocyanate.  Place  two  or  three  cm.3  of  a  KCN  solu- 
tion in  an  evaporating  dish  and  add  ammonium  polysulphide  (NH4)2SX 
until  the  color  of  the  latter  persists.  Evaporate  to  dryness.  Dissolve 
a  portion  of  the  residue  in  water;  add  a  few  drops  of  a  ferric  chloride 
(FeCl3)  solution  (?).  Test  a  solution  of  potassium  thiocyanate  (KCNS), 
shelf  reagent,  with  a  few  drops  of  FeCl3  ( ?).  If  the  solution  you  prepared 
gave  a  black  precipitate  when  FeCl3  was  added,  heat  the  residue  again 
in  the  evaporating  dish  and  repeat  test  with  FeCl3.  What  is  the  name 
of  the  prepared  substance  ?  Its  formula  ?  Equations  ?  What  is  the 
action  of  a  solution  of  potassium  thiocyanate  upon  solutions  of  ferrous 
salts  ? 

Test  the  tap  water  for  iron  by  evaporating  100  cm.3  to  small  bulk; 
add  a  few  drops  of  HNO3  (?)  and  heat  to  boiling.  Cool  the  solution, 
and  add  a  few  drops  of  a  KCNS  solution  (?). 

All  the  single  cyanides,  except  those  of  the  alkalies,  alkaline  earths 
and  mercury  (ic)  are  insoluble  in  water.  Barium  cyanide  is  diffi- 
cultly soluble. 

Experiment  XVII.— A  Study  of  Flames. 

(a)  Recall    or  repeat   the   experiment   with    the    Bunsen    burner.— 
Preliminary  Exercises. 

(b)  Light  a  Bunsen  burner.     Bring  a  cold  piece  of  brass  or  iron  wire 
gauze  down  upon  the  flame.     Observe  that  the  flame  does  not  pass 
through  the  gauze  (Fig.  42).    Is  this  due.  to  the  absence  of  combustible 
gases?    'Apply  a  lighted  match  above  the  gauze  (?).    Why  did  the  flame 
not  pass  through  the  gauze  ?     Turn  off  the  gas;  hold  a  piece  of  gauze  3  or 
4  cm.  above  the  top  of  the  burner;  turn  on  the  gas  and  hold  a  lighted 
match  above  the  gauze  (?).     Explain  results.     Define  "kindling  tem- 
perature."    What  is  the  principle  upon  which  the  miner's  safety  lamp 
(Fig  43)  is  constructed? 


CARBON   FAMILY.  239 

(c)  Open  and  close  the  holes  of  the  lighted  burner  several  times. 
Results?  Hold  a  piece  of  glass  tubing  in  the  "luminous  flame"  of  the 
burner  (?).  Open  the  holes  at  the  bottom  of  the  burner  and  hold  the 
blackened  rod  in  the  non-luminous  flame  ( ?).  Account  for  the  deposition 
of  carbon  and  its  disappearance.  What  is  the  probable  cause  of  the 
luminosity  of  the  flame?  Test  your  conclusion  as  follows:  Place  a 
small  quantity  of  powdered  charcoal  in  a  piece  of  glass  tubing  and  blow 
the  particles  of  carbon  into  one  of  the  holes  at  the  base  of  the  burner. 
What  is  the  effect  on  the  luminosity  of  the  flame  ?  Explain.  Why  do  a 
number  of  the  "burning  oils"  burn  with  a  "sooty"  flame?  Why  is  the 
hydrogen  flame  non-luminous?  What  is  the  source  of  the  powerful 
light  emitted  by  a  Welsbach  burner  ?  What  is  the  object  of  the  holes  at 
the  base  of  the  burner ?  Why  is  the  Bunsen  flame  non-luminous?  Why 


FIG.  42.  FIG.  43. — Davy's  Safety  Lamp. 

does  the  gas  ordinarily  burn  at  the  top  of  the  tube  and  not  on  the  inside  ? 
What  causes  the  flame  to  "strike  back"  occasionally  and  burn  inside 
the  tube? 

(d)  Slightly  bend  a  glass  tube  about  15  cm.  long;  introduce  the  shorter 
arm  in  the  flame  about  2  cm.  above  the  top  of  the  burner.     Light  the  gas 
issuing  from  the  tube.     What  are  your  conclusions  concerning  the  con- 
ditions existing  in  this  portion  of  the  flame  ?     Explain. 

(e)  Oxidizing  and  reducing  flames.     Examine  and  sketch  the  parts 
of  a  very  small  luminous  flame. 

Hold  a  piece  of  bright  copper  wire  horizontally  across  a  Bunsen  flame 
so  that  the  wire  cuts  the  inner  cone.  Observe  that  the  portion  of  the  wire 
in  the  inner  cone  remains  bright,  while  those  portions  in  contact  with  the 
edges  of  the  flame  become  coated  with  a  dark  substance  (copper  oxide, 
CuO).  The  copper  has  undergone  what  kind  of  a  chemical  change? 
Move  the  wire  so  as  to  bring  the  oxidized  portion  into  the  inner  cone, 
the  metal  becomes  bright  owing  to  a  reduction  of  the  oxide.  The  outer 


240 


EXPERIMENTAL  CHEMISTRY. 


cone  of  the  flame  where  oxygen  is  in  excess  is  called  the  oxidizing  flame ; 
the  inner  cone,  in  which  heated  and  unburnt  combustible  gases  exist 
(hydrogen  and  hydrocarbons),  is  referred  to  as  the  reducing 
flame.     Both  cones  exist  in  all  ordinary  flames. 

(/)  Use  of  the  blow-pipe.  Ask  the  instructor  how  to 
produce  the  oxidizing  and  reducing  flames  by  means  of  a 
blow-pipe.  Ascertain  the  effect  of  each  flame  by  heating  an 
intimate  mixture  of  lead  oxide  (PbO)  and  sodium  carbonate 
on  charcoal  in  the  reducing  cone,  and  a  small  piece  of 
metallic  lead  on  charcoal  in  the  oxidizing  cone.  Results  ? 

(g)  Examine  and  sketch  a  candle  flame  (Fig.  44).  Is  there 
any  essential  difference  between  a  gas  flame  and  a  candle 
flame  or  a  lamp  flame  ?  All  flames  are  the  result  of  the 
interaction  of  what  state  of  matter?  Define  "a  flame." 

(h)  Examine  a  lamp  burner.     Does  its  structure  embody 
principles  analogous  to  those  of  the  Bunsen  burner?     Ex- 
plain.    Why    are    some    lamps    provided    with    a    "central 
draft?"     Can  you  explain  the  cone-like  shape  of  the  flame 
FIG.  44.      of  t^  Bunsen  burner  ? 

The  following  table  gives  the  results  of  a  series  of  accurate  experiments 
conducted  for  the  purpose  of  determining  the  actual  temperatures  of 
various  regions  of  the  Bunsen  flame.  The  cooling  effect  of  the  diluent, 
air,  is  observed  in  portions  of  the  non-luminous  flame. 


TEMPERATURE    OF    FLAME    OF    BUNSEN    BURNER. 
(Data  for  a  burner  burning  six  cubic  feet  of  coal  gas  per  hour.) 


(Lewes-Newth.) 

Region  in  Flame.  Luminous. 

One-half  inch  above  burner     I35°-  •  •  • 

One  and  one-half  inch  above  burner  .  .  421°.  ... 

Tip  of  inner  cone 913  °.  .  .  . 

Center  of  outer  cone   1328°.  ... 

Tip  of  outer  cone 728°.  ... 

Side  of  outer  cone,  level  with  tip  of  in- 
ner cone 1236°. . . . 


Non-luminous. 
54° 


1090' 


1333' 


SILICON,  SI. 

At.  Wt.  28.4     Mol.  Wt.  (  ?  ). 

It  has  been  estimated  that  27.3  per  cent,  of  the  earth  consists  of  silicon. 
However,  owing  to  its  great  affinity  for  oxygen,  it  does  not  occur  in 
nature  as  an  uncombined  element.  In  the  combined  condition  it  most 
frequently  occurs  as  silicon  dioxide  (silica,  SiO2)  and  in  the  form  of  salts 
of  silicic  acid — silicates.  Silicon,  like  carbon,  occurs  in  several  forms, 


CARBON    FAMILY.  241 

of  which  amorphous  and  crystalline  forms  are  accurately  known. 
Amorphous  silicon  is  a  non-lustrous  greenish-brown  powder  which  when 
set  on  fire  in  the  air  undergoes  incomplete  combustion  owing  to  the 
formation  of  an  incombustible  layer  of  silicon  dioxide. 

Crystalline  silicon  (sp.  gr.,  2.49)  is  a  grayish-black  octohedral  crystal 
possessing  a  metallic  luster.  It  does  not  change  appreciably  in  the  air 
even  at  red  heat.  It  is  soluble  in  hot  caustic  soda.  (Molten  iron  or 
zinc  dissolves  silicon;  when  they  cool,  the  silicon  separates  in  crystals.) 

The  most  important  compound  of  silicon  is  silicon  dioxide,  the  anhy- 
dride of  silicic  acid.  It  occurs  in  several  v  arieties — two  crystalline  and 
one  amorphous.  It  is  most  widely  distributed  in  the  crystalline  form  as 
quartz  or  rock  crystal.  There  are  several  colored  varieties  of  quartz, 
the  color  being  due  to  the  presence  of  impurities.  Quartz  has  a  density 
of  2.66  and  stands  seventh  in  the  scale  of  hardness.  (See  Appendix.) 
The  other  crystalline  form  of  silicon  dioxide  is  known  as  tridymite.  It 
occurs  almost  solely  in  the  form  of  microscopic  crystals  as  a  constituent 
of  various  rocks.  The  sp.  gr.  of  tridymite  is  2.3.  Amorphous  silicon 
dioxide  (sp.  gr.,  2.2)  occurs  most  abundantly  in  the  form  commonly 
known  as  flint.  The  opal  represents  another  impure  form.  Jasper, 
chalcedony,  etc.,  are  "cryp to-crystalline,"  i.e.,  are  composed  of  small 
crystals.  Silica  dissolves  in  the  fused  hydroxides  and  carbonates  of  the 
alkalies,  but  is  insoluble  in  water  and  all  acids  except  hydrofluoric  acid. 
At  the  high  temperature  of  the  oxyhydrogen  flame  silicon  dioxide  melts 
forming  a  viscous  glass-like  fluid  which  can  be  blown  into  any  desired 
shape.  On  account  of  its  high  melting-point  and  resistance  to  chemical 
action,  this  amorphous  "quartz-glass"  has  been  used  extensively  in 
recent  years  for  making  vessels  for  use  in  chemical  laboratories.  It 
has  a  very  small  coefficient  of  expansion  with  heat  and  can,  therefore, 
be  subjected  to  sudden  changes  of  temperature  without  cracking. 

In  its  chemical  relations  silicon  bears  a  close  resemblance  to  carbon. 
It  is  essentially  non-metallic  in  these  relations.  The  natural  silicates 
form  a  numerous  and  complex  class  of  minerals  which  may  be  regarded 
as  compounds  of  metallic  oxides  with  silicon  dioxide;  as  feldspar,  A12O3,- 
K2O,  6SiO2. 

Experiment  I. — Preparation  and  Properties  of  Silicon. 

(a)  Procure  a  piece  of  crystalline  silicon  from  the  instructor.  Ex- 
amine it  carefully.  Tabulate  its  observed  properties.  What  is  its  sp.  gr.  ? 
Enumerate  the  several  forms  in  which  silicon  occurs. 

(6)  Thoroughly  mix  in  a  mortar  i  gram  of  magnesium  powder  and  2  or 
3  grams  of  powdered  quartz  sand.  Place  the  mixture  in  a  wide  hard  glass 
tube;  heat  the  tube  to  bright  redness  in  the  flame  of  a  blast-lamp,  slowly 
rotating  the  tube.  After  a  few  minutes'  heating,  dissolve  out  the  soluble 
portions  of  the  fused  mass  with  water  and  HC1.  The  residue  consists 
of  amorphous  silicon  and  quartz  sand. 

(c)  Ascertain  the  action  of   the  following  reagents  upon  silicon  and 
tabulate  results:     HC1,  H2SO4,  HNO3,  and  a  KOH  solution. 
16 


242  EXPERIMENTAL  CHEMISTRY. 

Experiment  II. — Silicon  Dioxide.     Silicates.    Silicic  Acid.    Hydrogele. 

(a)  Examine  specimens  of  quartz,  sand,  agate,  opal,  amethyst,  onyx, 
flint,    infusorial    or    diatomaceous   earth,    etc.      Make  a  list    of    those 
examined. 

(b)  Test  the   solubility  of  silicon  dioxide  in  the  various  acids  and 
alkalies.     Tabulate  results.      Recall  the  action  of  H2F2  on  silica  (SiO2). 
See  Exp.  Ill,  "Fluorine."     Equation  ? 

(c)  To  2  cm.3  of  a  solution  of  sodium  silicate  (water  glass,  Na2SiO3) 
in  an  evaporating  dish  or  casserole,  add  1 5  or  20  cm.3  of  water,  then  add  con- 
centrated HC1  until  the  solution  is  strongly  acid.    Note  the  formation  of  a 
white  gelatinous  precipitate.  What  is  it  ?   Formula  ?    Equation  ?    Note. — 
These  jelly-like  precipitates  are  sometimes  termed  "  hydro geles." 

Evaporate  the  solution  to  dryness  on  a  sand  or  steam  bath.  Now 
heat  the  dish  with  the  flame.  When  the  dish  has  cooled,  treat  the  residue 
with  water;  filter.  Remove  the  residue  from  the  filter  paper  and  heat  it 
in  a  covered  crucible  supported  by  a  pipe-stem  triangle.  Results  ?  What 
is  the  substance  in  the  crucible  ?  Its  formula  ?  Rub  some  of  it  between 
the  fingers.  Do  you  detect  any  grit  ?  Is  it  crystalline  or  amorphous  ? 

Sil-ic  acid  (probably  H4SiO4)  shows  a  tendency  to  form  "colloidal 
solutions"  or  "pseudo-solutions."  When  hydrochloric  acid  is  added 
to  a  concentrated  solution  of  a  soluble  silicate,  the  silicic  acid  separates 
out  as  a  gelatinous  mass,  but  if  a  dilute  solution  of  the  silicate  is  sub- 
stituted and  an  excess  of  the  acid  used,  precipitation  does  not  take  place 
and  the  solution  remains  clear  and  apparently  unchanged.  It  appears 
as  if  the  silicic  acid  were  difficultly  soluble  in  water,  but  dissolved  when 
much  water  was  used.  It  has  been  found,  however,  that  this  is  not  the 
case.  The  silicic  acid  is  present  (suspended)  in  the  water  in  the  colloidal 
state.  The  apparent  solution  of  silicic  acid  is  not  a  true  solution,  but  a 
"pseudo-solution."  (A  solution  of  this  kind  is  sometimes  termed  a 
hydrosole.)  This  may  be  proven  by  subjecting  the  pseudo-solution  to 
dialysis,  when  the  salt  (crystalloid)  formed  by  the  interaction  of  the 
metal  and  acid  will  diffuse  freely  through  the  membrane  of  a  dialyser, 
while  the  silicic  acid  will  be  retained,  as  is  in  accord  with  the  conduct 
of  colloidal  substances.  Again,  it  is  possible  to  partially  precipitate  the 
silicic  acid  from  these  "pseudo-solutions"  by  the  addition  of  various 
substances,  such  as  salts.  Silic  acid  occurs  very  frequently  in  this  form 
in  nature  as  the  result  of  the  decomposition  of  certain  silicates  by  carbonic 
acid.  The  various  colored  varieties  of  quartz  were  probably  formed 
from  silicic  acid  in  this  condition.  This  view  is  supported  by  the  fact 
that  it  is  now  known  that  silicic  acid,  under  suitable  conditions,  will 
crystallize  out  from  such  solutions. 

(b)  Fuse  a  mixture  of  i  grm.  of  finely  powdered  feldspar  with  4  or  5 
grams  of  sodium  carbonate*  in  a  covered  porcelain  or  platinum  crucible. 

*  A  mixture  of  K2CO3  and  Na2CO3  may  be  used  as  a  flux.  It  is  frequently  expe- 
dient to  use  a  fusion  mixture  consisting  of  i  part  of  NH4C1  and  8  parts  of  CaCO3  by 
weight. 


CARBON   FAMILY.  243 

The  crucible  should  not  be  more  than  half-filled.  Heat  gently  until 
the  frothing  has  ceased,  then  heat  with  the  blast-lamp  flame  until  the 
decomposition  is  complete  and  the  contents  of  the  crucible  are  in  quiet 
fusion.  When  the  crucible  is  cold  it  is  placed  upon  its  side  in  a  beaker 
containing  sufficient  water  to  cover  the  crucible.  Hydrochloric  acid  is 
now  added  until  effervescence  ceases,  and  no  further  precipitation  of 
gelatinous  silicic  acid  takes  place.  Filter.  Transfer  the  hydrogele  of 
silicic  acid  from  the  filter  paper  to  a  crucible  and  dehydrate  it.  Results  ? 
What  is  the  substance  which  remains  in  the  crucible  after  heating? 
Rub  some  of  it  between  the  fingers.  Do  you  detect  any  grit  ?  Equations? 

Experiment  IV. — Fluosilicic  Acid.     Fluosilicates. 

(a)  To  5  cm. 3  of  a  solution  of   fluosilicic  acid  (H2SiF6)  add  a  small 
quantity   of   a   solution   of   potassium   nitrate    (KNO3).     Examine   the 
precipitate  of  potassium  fluosilicate. 

(b)  Prove  that  fluosilicates  are  decomposed  when  heated  with  H2SO4, 
and  that  silicon  tetrafluoride  is  evolved.    Recall  Exp.  Ill — "Fluorine." 
Record  procedure.     Equations? 

(c)  Repeat  (a),  using  a  solution  of  a  sodium  salt.     Results?     Equa- 
tions ? 

(d)  Repeat     (a),    using    a    solution    of    barium    chloride.     Results? 
Equations  ? 

(e)  Record  the  names   and  formulae   of  the    well-defined   hydrides, 
halides   and   acids  of   silicon.      What  is  carborundum?     Its  formula? 

Experiment  V. — A  Test  for  Silicates.     Silicon  Tetraflouride. 

(Hood.)  Place  2  or  3  cm.3  of  concentrated  R2SO4  in  a  platinum 
crucible  and  add  a  small  quantity  of  powdered  silicate  or  substance  to 
be  tested.  Warm  the  contents  of  the  crucible  gently;  allow  crucible  to 
stand  until  effervescence  ceases.  When  the  contents  of  the  crucible 
have  cooled  (Caution!),  add  i  cm.3  of  H2F2.  Place  the  cover  on  the 
crucible  in  such  a  manner  as  to  leave  a  small  opening  on  one  side.  Hold 
a  platinum  loop  containing  a  drop  of  distilled  water  at  this  opening; 
warm  the  crucible  gently  and  observe  whether  white  particles  of  silicic  acid 
appear  in  the  drop  of  water.  Evaporate  the  drop  to  dryness  on  the 
crucible  cover  or  a  piece  of  platinum  foil.  A  white  residue  which  is  not 
volatilized  at  a  high  temperature,  proves  that  silica  was  present  in  the 
original  substance.  Equations? 

All  silicates  save  those  of  the  alkali  metals  are  insoluble  in  water. 
They  are  decomposed  by  hydrofluoric  acid  with  the  formation  of  silicon 
tetrafluoride.  A  very  few  of  the  insoluble  silicates  are  decomposed  by 
acids  other  than  hydrofluoric.  The  insoluble  silicates  are  usually 
converted  into  a  soluble  form  by  fusion  with  the  carbonates  of  the  alkalies. 


244 


EXPERIMENTAL  CHEMISTRY. 


The  physical  and  chemical  relations  of  carbon  and  silicon  may  be  seen 
in  the  following  table: 


Physical  Properties. 
Atomic  weight, 
State  or  phase, 

Color, 

Specific  gravity, 

Specific  heat, 
Melting-point, 

Chemical  Properties. 
H-derivatives, 
Heat  of  formation, 
State  or  phase, 

Halides 

Heat  of  formation, 

State  or  phase, 

O-derivatives, 
Heat  of  formation, 
State  or  phase, 

S-derivatives, 
Heat  of  formation, 
State  or  phase, 


Carbon. 

12.0 

Solid 

(amorphous,  crystalline) 

Black  (amorph.) 
Black-gray  (graph.) 

2.2-3.5  (cryst.) 
(graphite,   diamond) 


Silicon. 
28.4 
Solid 

(amorphous,  crystalline) 

Greenish-brown    (amorph.) 

Grayish-black  (cryst.) 

2.4 
(crystalline) 


35°° 


CH4;  C2H4,  etc. 
C,  H4  =  2i,8oocal. 
Gas 

CC14;  etc. 

C,C14  =  21,000  cal. 

Liquid 

CO;  C02 

C,O2  =96,980  cal. 
Gas 

CS2 

C,  S2=  — 26,100  cal. 

Liquid 


SiH4 

SiH4  =  24,800  cal. 

Gas 


SiCl4;_ 
Liquid 


cal. 


SiO2 

Si,O2  =  2 19,000  cal. 

Solid 

SiS2 

Si, 82  =  40,400  cal. 
Solid. 


PROBLEMS. 

1.  Five   grams  of   pure   graphite  are  completely  burned  in  oxygen. 
What  volume  of  carbon  dioxide  is  formed?     What  is  the  heat  of  the 
reaction  ? 

2.  What  volume  of  oxygen  is  necessary  to  burn  12  grams  of  carbon? 

3.  How  many  grams  of  pure  calcium  carbonate  will  be  required  to 
produce  10  liters  of  carbon  dioxide? 

4.  What  volume  of  carbon  monoxide  will  be  liberated  by  the  action 
of  H2SO4  on  10  grams  of  oxalic  acid? 

5.  An  analysis  of  the  air  in  a  lecture-room  showed  that  8.5  volumes 
of  carbon  dioxide  were  present  in    10,000  volumes  of  air.     If  the  room 
is  20  meters  long,  17  meters  wide  and  7  meters  high,  what  is  the  weight 
and  volume  of  the  carbon  dioxide  in  the  room  ? 

6.  Dumas  and   Stas  found  that  30  parts  of  carbon  by  weight  com- 
bined with  So  parts  of  oxygen,  and  that  the  carbon  dioxide  formed  con- 
tained its  own  volume  of  oxygen.     A  liter  of  carbon  dioxide  weighs 
1.976  grams.     From  the  foregoing  data  calculate  the  molecular  weight 
and  deduce  the  simplest  formula  of  carbon  dioxi  de.     Deduce  the  atomic 
weight  of  carbon. 

7.  How  many  liters  of  carbon  dioxide  must  be  passed  over  red-hot 


CARBON    FAMILY.  245 

charcoal  to  yield  100  liters  of  carbon  monoxide?     To  yield  25  grams  of 
carbon  monoxide? 

8.  A  precipitate  which  when  dry  weighed  1.5  grams  was  formed  by 
passing  carbon  dioxide  through  lime-water.     What  is  the  weight  of  the 
absorbed  gas? 

9.  Calculate    the   precentage  weight   of  CO2  in  calcium  carbonate, 
sodium  bicarbonate,  and  magnesium  carbonate. 

10.  How  many  grams  of  oxygen  will  be   required  for  the  complete 
combustion  of  5  grams  of  methane  ?     Of  5  grams  of  ethylene  ?     Of  5 
grams  of  acetylene  ? 

11.  Calculate  the  percentage    composition    of    marsh  gas,   ethylene 
and  acetylene. 

12.  Water  gas  is  prepared  by  passing  steam  over  red-hot  coal.     The 
simplest  equation  for  the  reaction  is: 

C  +  H2O  —  H2  +  CO. 

What  volume  of  water  gas  at  15°  C.  and  730  mm.  will  be  formed  from 
50  grams  of  steam? 

13.  Scheele   found   that  0.6738  of   silicon  tetrachloride  reacted  with 
AgNO3  to  yield  2.2769  grams  of  silver  chloride.     Calculate  the  atomic 
weight  of  silicon. 

14.  Calculate  the  percentage  composition  of 

Garnet,  Ca3Fe2  •  •  •  (SiO4)3 

Mica,  KH2Al3(Si04)3 

Kaolin,  H2Al2(SiO4)3,  H2O 

Serpentine,        Mg3Si2O7,2H2O 

Orthoclase  (feldspar),  KAlSi3O8. 


CHAPTER  XXI. 


SOME  COMMON  CARBON  COMPOUNDS. 

The  simplest  compounds  of  carbon  are  those  which  contain  only  hydro- 
gen and  carbon,  e.g.,  CH4,  C6H6,  C2H4.  These  compounds  have  been 
termed  hydrocarbons  and  are  regarded  as  the  fundamental  compounds 
of  organic  chemistry.  Nearly  all  organic  compounds  are  either  hydro- 
carbons or  hydrocarbon  derivatives.  Speaking  comparatively,  very  few 
carbon  compounds  are  known  which  do  not  contain  hydrogen. 

Although  hundreds  of  hydrocarbons  are  known  and  an  almost  in- 
finite number  are  theoretically  possible,  they  may  be  arranged  in  a 
small  number  of  comparatively  simple  series.  These  series  correspond 
somewhat  to  the  different  groups  of  elements.  Further,  the  members 
of  one  and  the  same  series  resemble  one  another  much  more  closely  than 
do  the  elements  of  a  given  series  of  the  elements.  In  general,  the  members 
of  any  series  bear  such  a  close  resemblance  to  one  another,  that  if  we 
understand  the  simpler  members,  we  are  able  to  anticipate  many  of  the 
properties  of  the  more  complicated  members.  Attention  is  also  directed 
to  the  fact  that  for  each  hydrocarbon  in  a  series  there  is  a  corresponding 
class  of  derivatives  and  that  the  relations  existing  between  any  hydro- 
carbon and  its  derivatives  are  very  similar  to  those  existing  between  any 
other  hydrocarbon  of  the  series  and  its  derivatives.  It  is  obvious,  then, 
that  if  we  know  the  derivatives  which  can  be  yielded  by  a  single  hydro- 
carbon of  a  series,  then  we  are  able  to  prophesy  with  some  degree  of 
certainty  the  existence  of  the  derivatives  of  every  other  hydrocarbon  of 
the  series. 

The  following  table  shows  the  classification  of  a  number  of  the  hydro- 
carbons as  regards  their  empirical  formulae  in  three  of  the  best  known 
"series."  It  will  be  observed  that  each  hydrocarbon  differs  from  the 
one  which  precedes  it  by  an  atom  of  carbon  and  two  of  hydrogen.  Such 
a  series  is  known  as  an  homologous  series  and  may  be  represented  by  one 
general  formula,  Cn  H2n+2. 


Methane  Series. 


Ethylene  Series. 
CH 


Methane  .....    C  H4 

Ethane  ......    C2H6 

Propane  ......    C3H8 

Butane  ......    C4H10 

Pentane 
Hexane 
Etc. 


C5H12 


Ethylene 
Propylene  . 
Butylene 
Amylene  .. 
Hexylene.. 
Heptylene  . 

246 


C2H4 

C3H6 

C4H8 

C5H10 

C6H12 

C7H14 


Benzene  Series. 
CnH2n_6. 

Benzene C6H6 

Toluene C7H8 

Xylene C8H10 

Mesitylene  .  .    C9H12 
Pseudocumene  C9H12 

Durene CIO^M 

Cymene C10H14 


SOME  COMMON  CARBON  COMPOUNDS  247 

Saturation.  —  The  hydrocarbons  are  frequently  referred  to  as  being 
saturated  or  unsaturated  accordingly  as  they  behave  towards  chemcal 
agents.  Thus,  if  all  the  four  valences  of  carbon  are  employed,  the 
hydrocarbon  having  then  no  power  to  combine  directly  with  other  com- 
pounds or  elements,  it  is  said  to  be  saturated. 

CH4  +  C12  ->  CH3C1  +  HC1. 

It  is  obvious  that  in  the  reaction  indicated  above  chlorine  must  first 
displace  hydrogen  before  it  can  enter  into  combination  with  the  com- 
pound. On  the  other  hand,  if  the  hydrocarbon  (any  compound)  can 
unite  directly  with  elements  or  compounds  it  is  spoken  of  as  being  un- 
saturated; for  example,  ethylene  combines  directly  with  chlorine  to 
form  ethylene  chloride  — 

C2H4  +  C12->C2H4C12. 

Organic  Radicals.  —  The  hydrocarbons  may  be  regarded  as  hydrides  of 
various  radicals,  as  methyl  hydride  or  methane,  CH3.H,  ethyl  hydride 
or  ethane,  C2H5.H,  etc.  As  a  matter  of  fact,  hydrocarbons  from  which 
hydrogen  has  been  removed  give  rise  to  hydrocarbon  radicals,  thus, 
CH3  is  the  organic  radical  methyl  from  methane,  and  C2H5  is  the  radical 
ethyl  from  ethane.  These  groups  of  units,  radicals,  which  are  found 
to  pass  unaltered  from  compound  to  compound  are  like  all  other  radicals, 
incapable  of  existing  in  the  jree  state.  Since  in  chemical  reactions 
these  organic  radicals  behave  toward  other  elements  and  radicals 
in  a  manner  similar  to  that  in  which  the  metals  behave  toward  the  non- 
metals  they  are  frequently  termed  positive  radicals.  They  usually 
differ  from  the  inorganic  radicals  in  that  they  lack  the  property  of  form- 
ing ions.  In  general,  organic  substances  may  be  most  conveniently 
regarded  as  derived  from  these  hyrdocarbon  radicals,  which  in  most 
cases  are  hydrocarbons  possessing  unsatisfied  bonds  of  valency. 

The  following  scheme,  containing  a  partial  list  of  the  more  important 
fundamental  forms  of  the  hydrocarbon  derivatives  may  be  of  service 
in  emphasizing  the  thought  of  the  foregoing  paragraph: 

Hydrocarbons   ................  R  —  H 

Alcohols  .....................  R  —  O  —  H 

Phenols  ......................  R  —  O  —  H 

Ethers  .......................  R—  O  —  R 

Aldehydes    ...................   R  —  C 

Acids    .......................   R  —  C      o—  H 

Ketones    .....................   R  —  C  —  R 


Acid  anhydrides  ..............   R  —  C  /C—  R 

Halides  .  .   RC1,  RBr,  etc. 


248  EXPERIMENTAL  CHEMISTRY. 

Esters R  —  C  \Q_ R 

Amides     R""C\NH2 

Amines     RNH2,  R2NH2,  etc. 

Organo-Mineral  Compounds  ...  R  —  M 

Carbohydrates C6H10O5,  C12H22O11,  etc. 

Explanatory  Note.* 

Alcohols. — Composed  of  carbon,  hydrogen  and  oxygen  and  containing 
one  or  more  hydroxyl  groups  (OH),  as  methyl-alcohol,  CH3OH,  ethyl- 
alcohol,  C2H5OH. 

Phenols. — Similar  to  alcohols  in  composition,  but  resembling  the 
acids  in  many  of  their  properties;  however,  they  do  not  yield  aldehydes 
when  partially  oxidized  as  phenol,  C6H5 .  OH. 

Ethers. — The  oxides  of  the  radicals;  formed  from  the  alcohols  by 
the  substitution  of  a  hydrocarbon  radical  for  the  H  in  the  hydroxyl, 
as  methyl-ether,  CH3 .  O .  CH3 ;  ethyl-ether,  C2H5 . 0 .  C2H5. 

Aldehydes. — Dehydrogenated  alcohols;  products  of  the  partial  oxida- 
tion of  the  alcohols,  containing  the  group  (COH),  as  formaldehyde,  CH2O; 
as  ethyl-aldehyde  (acet-aldehyde)  CH3.COH. 

Acids. — Products  of  the  further  oxidation  of  the  alcohols,  containing 
one  or  more  carboxyl  radicals,  CO2H;  as  formic  acid,  CH2O2;  acetic  acid, 
CH3  CO2H  or  CH3 .  CO .  OH. 

Ketones. — (Acetone,  the  simplest  of  the  ketones,  is  prepared  by  the 
distillation  of  calcium  acetate.)  Formed  from  the  acids  by  the  substi- 
tution of  a  hydrocarbon  radical  for  the  OH  in  the  carboxyl;  contain  the 
group  CO,  as  acetic  tetone  or  acetone  CH3.  CO .  CH3. 

Halides. — Formed  from  the  foregoing  groups  by  the  substitution  of  a 
halogen  radical  for  hydrogen  or  hydroxyl;  as  chloroform,  CHC13;  iodo- 
jorm,  CHI3,  ethyl  chloride,  C2H5C1;  acetyl  chloride,^  CH3 .  CO .  Cl. 

Esters  (ethereal  salts). — Formed  from  the  acids  by  the  substitution 
of  a  hydrocarbon  radical  for  the  hydrogen  in  the  carboxyl  radical;  as 
ethyl  acetate,  CH3 .  CO .  O .  C2H5. 

Ammonia — Derivatives. — Formed  upon  the  model  of  ammonia,  NH3, 
by  the  substitution  of  a  radical  for  hydrogen;  as  acetamide,  NH2.C2H3O; 
ethylamine,  NH2.C2H5. 

Organo-Mineral  Compounds. — Formed  upon  the  type  of  the  chlorides 
of  metals  or  non-metals  by  the  substitution  of  hydrocarbon  radicals 
for  the  chlorine,  as  zinc  ethide,  Zn(C2H5)2. 

Compounds  for  which  the  exact  structure  has  not  been  fully  estab- 
lished. They  are  classified  according  to  similarity  of  properties,  ultimate 
composition  or  products  of  decomposition. 

Carbohydrates,  or  compounds  usually  containing  six,  or  some  multiple 
of  six,  atoms  of  carbon,  together  with  some  multiple  of  the  group  H2O,f 
as  starch,  C6H10O5;  glucose,  C6H12O6;  sugar,  C12H22On. 

*Data  from  Bloxam.  fRhamose  (C6HI2OS)  is  an  exception. 


SOME    COMMON    CARBON   COMPOUNDS.  249 

Glucosides,  or  compounds  which  yield  glucose  as  one  of  their  products 
of  decomposition,  as  salicin,  C13H18O7. 

Albuminoids  and  Gelatinoids,  or  compounds  containing  C,  H,  N  and 
O,  often  with  small  quantities  of  S,  and  sometimes  of  P,  distinguished 
by  their  tendency  to  putrefy  when  moist;  albumin,  fibrin  and  casein 
are  examples  of  such  compounds,  but  they  cannot  at  present  be  repre- 
sented by  satisfactory  formulae. 

Isomerism. — One  of  the  most  interesting  facts  with  which  we  early  be- 
come acquainted  in  the  study  of  organic  chemistry  is  the  existence  of 
entirely  different  compounds  having  the  same  percentage  composition. 
Such  compounds  are  called  isomeric  or  briefly  isomers.  Isomers  are 
of  two  kinds:  (i)  If  the  isomers  have  the  same  molecular  weight  they 
are  known  as  metamers,  as  ethyl  alcohol  (C2H5OH)  and  methyl-oxide 
(  (CH3)2O).  (2)  If  they  have  different  molecular  weights,  as  acetylene 
(C2H2),  benzene  (C6H6)  and  styrene  (C8H8),  they  are  called  polymers. 

Isomers  usually  differ  in  both  their  physical  and  chemical  properties. 
For  example,  the  empirical  formula  C2H6O  represents  two  substances 
(metamers) — the  one  a  gas  (methyl-oxide),  the  other  a  liquid  (ethyl 
hydroxide  or  alcohol).  Isomeric  compounds  undoubtedly  owe  their 
differences  to  different  groupings  or  arrangements  of  the  atoms  within 
the  molecules.  It  frequently  occurs  that  one  of  the  metamers  can  be 
made  directly  from  the  other.  This  gave  rise  to  that  variety  of  chemical 
change  previously  named  "internal  rearrangement."*  In  1828,  Wohler 
effected  the  synthesis  of  a  well-defined  organic  compound,  urea,  CO- 
(NH2)2,  from  wholly  inorganic  materials.  This  was  the  first  synthesis 
by  a  chemist,  of  a  typical  organic  compound  from  a  substance  whose 
preparation  is  independent  of  life  processes.  Wohler  warmed  a  solution 
of  ammonium  cyanate,  NH4.CNO,  which  is  a  metamer  of  urea,  for  some 
time.  On  cooling  the  liquid,  long  prisms  of  urea  were  deposited. 

NH4CNO  +±  CO(NH2)2. 

These  two  compounds  are  entirely  different  as  regards  their  chemical 
properties.  Ammonium  cyanate  is  an  ionizable  salt  while  the  latter 
(urea)  is  not  a  salt,  but  unites  with  acids  like  ammonia  to  form  salts. 
The  ammonium  cyanate  undoubtedly  underwent  an  intramolecular 
change  as  indicated  by  above  formulae;  at  least  the  attempt  is  made  to 
explain  the  difference  in  properties  of  the  metamers  by  postulating  a 
difference  in  the  molecular  structure  or  constitution  oj  the  substances. 

It  is  quite  probable  that  a  difference  in  structure  is  closely  related 
to  a  difference  in  energy-content  of  metamers.  For  example,  fumaric 
and  malei'c  acids  have  the  same  empirical  formula,  C4H4O4,  but  the 
"heat  of  combustion"  of  the  former  is  1338 Kj.,  while  that  of  malei'c 
acid  is  1365  Kj.  This  indicates  that  the  energy-content  of  malei'c  acid 
is  greater  than  that  of  fumaric  acid.  The  conclusion  seems  justified 

*  One  author  suggests  that  there  may  be  discovered  802  compounds  of  the 
ormula,  d3H28. 


250  EXPERIMENTAL  CHEMISTRY. 

when  it  is  recalled  that  maleiic  acid  is  the  less  stable  of  the  two  compounds 
and  melts  at  130°,  while  fumaric  acid  melts  at  a  temperature  above  200°. 
The  following  structural  formulae  indicate  a  difference  in  the  constitution 
of  the  two  compounds: 

COOH— C— H  CH— COOH 

II  II 

H  —  C  —  COOH  CH— COOH 

Fumaric  acid  Male'ic  acid 

(C4H404)  (C4H404) 

Graphic  or  structural  formulae  are  the  result  of  an  endeavor  to  provide 
a  "picture"  of  the  way  in  which  the  atoms  or  radicals  of  a  molecule  are 
linked  together. 

It  is  acknowledged  that  any  success  which  has  been  or  may  be  achieved 
in  the  synthesis  of  organic  compounds  has  been  or  will  be  the  result  of 
the  study  of  the  structure  of  compounds. 

The  proving  of  the  constitution  of  a  substance  represents  in  many 
cases  one  of  the  most  difficult  problems  of  organic  chemistry.  It  usually 
involves  an  accurate  and  extended  consideration  of  the  reactions  of  the 
substance  under  consideration.  In  order  that  the  student  of  inorganic 
chemistry  may  have  some  conception  of  the  mode  of  procedure  of  a 
chemist  endeavoring  to  prove  the  constitution  of  a  substance,  the  follow- 
ing method  is  cited,  and  may  be  regarded  as  typical  although  it  represents 
one  of  the  simpler  cases. 

Determination  oj  the  Constitution  oj  Alcohol. — (a)  When  alcohol 
reacts  with  hydrochloric  acid,  and  the  mixture  is  distilled  at  a  low  tem- 
perature, the  products  are  water  and  a  volatile  liquid  known  as  ethyl 
chloride,  C2H5C1.  The  interaction  may  be  indicated  by  the  equation, 
C2H5OH  +  HC1  ->  C2H5C1  +  HOH.  It  is  evident  that  the  Cl  of  HC1 
has  exchanged  positions  with  the  OH  of  the  alcohol;  this  leads  to  the 
conclusion  that  alcohol  is  composed  of  at  least  two  groups,  namely  the 
ethyl  radical  and  the  hydroxyl  radical,  and  that  its  rational  formula  is 
C2H5.OH. 

(b)  Again,  sodium  dissolves  in  alcohol  (C2H5OH)  with  an  evolution 
of  hydrogen,  forming  a  crystalline  substance  known  as  sodium  ethoxide, 
C2H5ONa.  It  is  seen  that  Na  has  displaced  one  atom  of  H.  It  might 
be  supposed  if  an  excess  of  Na  were  used,  that  the  remaining  H5  might 
be  displaced  and  a  compound  of  the  composition  of  C2Na6O,  would  be 
produced  ultimately,  but  this  is  contrary  to  all  experimental  evidence. 
Na  can  be  substituted  for  only  one  of  the  six  atoms  of  hydrogen  in 
alcohol.  It  is  inferred  then  that  one  of  the  six  atoms  is  "linked"  or 
"fitted"  into  the  intramolecular  structure  in  a  manner  different  from  the 
other  five.  The  concept  is  indicated  by  writing  the  formula  for  alcohol 
as  follows,  C2H5OH. 

Chemists  have  been  unable  as  yet  to  agree  on  the  rational  formula 
of  many  compounds,  especially  the  carbohydrates,  glucosides,  alkaloids, 


SOME  COMMON  CARBON  COMPOUNDS.  251 

etc.     The  following  structural  formula  for  cane-sugar  (C12H22OU)  has 
been  proposed: 

CH2OH 
I 
CH2OH-CHOH.CHCHOH.CHOH-CH-O-C-CHOH.CHOH.CH-CH2OH). 

I                             II  ! 

|  o        o | 

Experiment  I. — Hydrocarbons. 

(a)  Recall  or  repeat  the  experiments  in  which  methane,  ethylene  and 
acetylene  were  prepared.  Are  these  hydrocarbons  gases,  liquids  or 
solids?  Write  the  equations  for  the  combustion  of  these  three  com- 
pounds. 

(6)  Examine  small  quantities  of  each  of  the  following  hydrocarbons 
and  give  the  formula  (of  the  main  components  if  a  mixture)  for  each. 
Classify  them  as  to  state:  Gasoline,  toluol,  kerosene,  naptha,  benzine, 
petroleum,  paraffine,  benzene,  ozokerite,  vaseline  and  gasoline. 

Experiment  II.— (L.  T.)  Determination  of  the  "Flashing-point"  of 
Kerosene. 

Note. — "The  temperature  at  which  oil  gives  off  sufficient  vapor  to 
form  a  momentary  flash  when  a  small  flame  is  brought  near  its  surface" 
is  known  as  the  "flash-point."  The  temperature  at  which  oil  gives 
off  enough  vapor  to  maintain  a  continuous  flame  if  ignited,  is  ascer- 
tained by  the  "fire  test."  The  " burning- point"  is  about  10°  C.  higher 
than  the  "  flash-point."  The  safety  of  kerosene  depends  largely  on  the 
absence  of  volatile  hydrocarbons  which  may  escape  in  sufficient  quan- 
tities to  form  an  explosive  mixture  with  air.  The  presence  of  these 
hydrocarbons  is  determined  by  the  flashing  test. 

(a)  The  simplest  form  of  apparatus  used  for  determining  the  flash- 
point consists  essentially  of  a  small  beaker  in  which  is  suspended  a 
thermometer.    A  small  quantity  of  the  kerosene  to  be  tested  is  placed  in  the 
beaker  and  warmed  slowly  till  a  flame  which  is  brought  near  to  the 
surface  at  regular  intervals  of  time,  causes  a  momentary  flash.     At  the 
flashing-point  the  vapor  ignites,  and  the  bluish  flame  runs  down  to  the 
surface  of  the  oil. 

(b)  A  simple,  but  more  accurate  form  of  apparatus  for  determining 
the  flashing-point  may  be  prepared  as  follows:     A  rubber  stopper  pro- 
vided with  single  perforation  is  fitted  to  one  end  of  a  piece  of  a  glass 
cylinder  2.5  cm.  in  diameter,  and  15  cm.  long.     One  end  of  a  piece 
of  ordinary  glass  tubing  about  40  cm.  long  is  forced  through  the  cork 
until  it  projects  about  i  cm.  beyond  the  inner  surface  of  the  cork;  the 
tubing  is  then  bent  off,  close  to  the  cork,  twice  at  right    angles  in  such 
a  manner  that  the  long  end  of  the  tube  is  parallel  and  close  to  the  cylinder; 
the  tubing  is  finally  bent  on  a  level  with  the  open  end  of  the  cylinder  at 
right  angles  and  away  from  the  latter.     The  apparatus  if  made  according 


252  EXPERIMENTAL  CHEMISTRY. 

to  directions,  has  much  the  appearance  of  the  old-fashioned,  long-stemmed 
pipes.  The  bent  tube  contracts  to  a  small  orifice  within  the  cork.  Air 
is  forced  through  this  tube.  The  oil  to  be  tested  is  poured  into  the 
cylinder  until  the  latter  is  filled  to  a  point  such  that  when  air  is  being 
forced  through  the  apparatus,  the  surface  of  the  foam  is  about  5  cm. 
from  the  top  of  the  cylinder.  The  apparatus  is  now  placed  in  a  beaker 
of  water — the  surface  of  the  oil  and  the  water  should  be  at  the  same 
level.  Suspend  a  thermometer  in  the  oil.  Heat  the  contents  of  the 
beaker  gently  and  force  a  slow  current  of  air  through  the  oil.  Bring  a 
small  flame  to  the  mouth  of  the  cylinder  for  an  instant  as  the  temperature 
of  the  kerosene  rises  slowly,  degree  by  degree.  The  lowest  temperature 
at  which  the  vapor  ignites,  as  indicated  by  the  bluish  flame  running  down 
to  the  surface  of  the  oil,  is  the  flash-point. 

Determine  the  flashing-points  of  two  or  three  grades  of  kerosene. 

From  what  you  know  of  the  properties  of  gasoline,  would  you  infer 
that  its  flash-point  is  higher  or  lower  than  that  of  kerosene  ? 

Experiment  III. — Fermentation.  Properties  and  Preparation  of  Al- 
cohol. 

(a)  Dissolve  50  grains  of  grape  sugar  (C6H12O6)*  in  350  cm.3  of  water, 
and  add  one-half  of  a  compressed  yeast  cake.  Place  the  mixture  in  a 
large  flask  (Fig.  45)  or  bottle  provided  with  a  one-hole  rubber  stopper; 


FIG. 


45- 


connect  the  flask  with  a  wash  bottle  half  filled  with  lime-water.  The 
delivery  tube  of  the  wash  bottle  should  be  connected  in  series  with  a 
U-tube  filled  with  caustic  potash  or  soda-lime,  the  object  of  which  is  to 
prevent  carbon  dioxide  in  the  air  from  acting  upon  the  lime-water.  Set 
the  apparatus  aside  in  a  moderately  warm  place.  Fermentation  usually 
begins  at  once  as  is  evidenced  by  the  bubbling  of  carbon  dioxide,  one  of 

*  Cane  sugar  (ordinary  sugar)  does  not  ferment.     If  boiled  with  acid  it  is  inverted, 
as  indicated  by  the  following  equation: 

CMHM01X  +  H20  —  C6HI206  +  C6HI306 
when  fermentation  may  take  place. 


SOME  COMMON  CARBON  COMPOUNDS.  253 

the  products,  through  the  lime-water.  This  action  should  be  allowed 
to  continue  for  two  days,  then  decant  the  liquid  upon  a  filter.  Set  up  a 
condenser  and  a  distilling  flask;  place  the  nitrate  together  with  several 
"boiling  tubes"  in  the  distilling  flask;  insert  a  thermometer  in  a  hole  in 
the  cork  with  which  the  distilling  flask  is  provided,  so  that  the  bulb  just 
touches  the  surface  of  the  contents  of  the  flask;  heat  gently,  and  collect 
that  portion  of  the  distillate  which  passes  over  between  80°  and  93°  C. 

(b)  Notice  the  odor  of  the  distillate  (?).     Place  a  few  drops  of  it  on  a 
watch  glass  and  ascertain  whether  it  burns  (?).     Test  its  reaction  with 
litmus  paper  (?). 

(c)  Repeat  (b),  using  (a)  ethyl  alcohol,  (b)  methyl  alcohol  (?).     What 
are  your  conclusions  as  to  the  nature  of  the  distillate  used  in  (b)  ?     What 
is  its  empirical  formula  ?     Its  rational  formula  ? 

(d)  Try  the  solubility  of  iodine,  camphor,  rosin,  shellac,  etc.,  in  sepa- 
rate portions  of  alcohol  (?). 

C6H12O6  —  2C2H5OH  +  2CO2. 

Ethyl  alcohol,  C2H5OH,  is  the  best-known  member  of  the  class  of  alco- 
hols. When  pure,  it  is  a  transparent,  colorless,  mobile  and  volatile  liquid 
possessing  a  peculiar  and  agreeable  odor.  It  has  a  sp.  gr.  of  0.797,  boils 
at  78.3°  C.,  solidifies  at — 130°  C.,  is  of  a  neutral  reaction,  and  burns  with 
a  non-luminous  flame.  When  mixed  with  water  a  contraction  of  volume 
occurs  and  heat  is  liberated.  The  attraction  of  alcohol  and  water  for 
one  another  is  so  great,  that  strong  alcohol  abstracts  it  from  tissues, 
membranes,  and  other  similar  substances  immersed  in  it;  to  this  property 
are  due  its  coagulating  action  on  albumin  and  its  preservative  action  on 
animal  substances.  Pure  alcohol  taken  internally  acts  poisonously; 
when  taken  in  the  dilute  form  it  possesses  intoxicating  properties.  It  is 
affirmed  by  various  authorities  that  "it  lowers  the  temperature  of  the 
body  from  0.5  °  to  2  °  C.,  although  the  sensation  of  warmth  is  experienced." 

Alcohol  which  does  not  contain  more  than  i  per  cent,  of  water  is  known 
as  pure  or  absolute  alcohol.  Diluted  alcohol,  made  by  mixing  equal 
volumes  of  water  and  alcohol,  has  a  sp.  gr.  of  0.936  and  is  identical  with 
the  proof-spirit  of  the  United  States  Custom-house  and  Internal  Revenue 
service. 

Alcohol  is  an  excellent  solvent  for  many  organic  and  inorganic  sub- 
stances. These  alcoholic  solutions  are  given  such  names  as  spirits, 
tinctures,  fluid  extracts,  etc.  The  terminology  depends  on  various  phar- 
maceutical relationships.  "Spirits  from  the  pharmaceutical  point  of 
view,  are  simply  alcoholic  solutions  of  volatile  substances."  "The  active 
ingredient  may  be  solid,  liquid  or  gaseous."  "Tinctures  are  alcoholic 
solutions  of  medicinal  substances.  They  differ  from  spirits  in  being 
made  from  non-volatile  bodies — the  tincture  of  iodine  being  the  one 
officinal  exception  to  this  rule."  "Fluid-extracts  are  liquid  alcoholic 
preparations  of  uniform  and  definite  strength,  made  by  percolating 
drugs  with  menstrua,  and  concentrating  a  portion  of  the  percolate  so  that 
in  each  case  a  cubic  centimeter  represents  the  medicinal  virtues  of  one 


254  EXPERIMENTAL  CHEMISTRY. 

gram  of  the  drug;  they  are  mostly  concentrated  tinctures."  —  "Practice  of 
Pharmacy.  "—Remington. 

Experiment  IV.  —  Tests  for  Ethyl  Alcohol. 

(a)  lodoform  test.     Dissolve  a  small  flake  or  crystal  of  iodine  in  2 
or  3  cm.3  of  alcohol;  add  a  strong  solution  of  crystallized  sodium  car- 
bonate or  potassium  hydroxide  until  the  brown  color  of  the  solution  dis- 
appears; the  yellow  precipitate  which  forms  is  iodoform,  CHI3. 

Note.  —  The  student  is  reminded  that  other  alcohols,  acetone,  aldehyde, 
etc.,  show  the  same  reaction. 

(b)  Aldehyde  test.     Reducing  action  of   alcohol.     To  4  or  5   cm.  3 
of  a  potassium  dichromate  solution  add  i  cm.3  of  sulphuric  acid  and 
2  cm.3  of  alcohol;  heat  the  solution  gently  and  note  the  odor  of  the 
gaseous  product,  acetic  aldehyde   (C2H4O),  which  is  a  volatile  liquid 
boiling  at  20.8°  C.     The  green  color  of  the  solution  is  due  to  the  for- 
mation of  chromium  sulphate,  Cr2(SO4)3. 

(c)  Ether  test.     Add  to  2  cm.3  of  alcohol  the  same  volume  of  strong 
sulphuric  acid;  heat  to  boiling  and  add  a  little  more  alcohol,  drop  by 
drop.     The  odor  of  ethyl  ether,  (C2H5)2O,  may  be  easily  detected  on 
further  heating. 

2C2H5OH  +  H2SO4-*  C5>SO4  +  H2O. 


|5        S04  +  C2H5OH-^  p250  +  H2S04. 

n       /  ^2n5  / 

(d)  Ester  test.  To  4  or  5  cm.3  of  a  strong  solution  of  sodium  acetate, 
NaC2H3O2,  add  a  few  drops  of  concentrated  sulphuric  acid  and  5  or  6 
drops  of  ethyl  alcohol,  then  warm  the  mixture  gently.  Notice  the 
pleasant  fruit-like  odor  of  the  ester,  ethyl  acetate,  a  volatile  liquid  which 
is  formed  under  the  above  conditions.  The  foregoing  is  a  test  for  either 
alcohol  or  acetic  acid. 

C2H5OH  +  C2H3O2H  —  C2H5C2H3O2  +  H2O 
Alcohol.        Acetic  acid.  Ethyl  acetate. 

C2H5OH  +  NaC2H3O2  +  H2SO4—  C2H5C2H3O2  +  NaHSO4  +  H2O. 

In  many  of  their  relations  the  alcohols  are  analogous  to  metallic 
hydroxides.  For  example,  CH3OH  and  C2H5OH  are  compounds  anal- 
ogous to  NH4OH,  KOH,  NaOH,  etc.  Both  classes  of  compounds  react 
with  acids  to  form  salts.  Ethyl  acetate,  one  of  the  products  of  the  inter- 
action of  ethyl  alcohol  and  acetic  acid,  is  an  organic  salt  analogous  to 
ammonium  or  sodium  acetate.  These  organic  salts  are  often  called  esters 
or  ethereal  salts.  The  latter  term  is  unfortunate  as  these  salts  do  not 
ionize  to  any  extent  and  possess  practically  none  of  the  properties  of  or- 
dinary salts. 


SOME    COMMON    CARBON    COMPOUNDS.  255 

Experiment  V. — Properties  of  a  Trihydric  Alcohol  (Glycerine  or  Gly- 
cerol). 

Place  a  little  glycerine  in  a  test  tube.  Note  its  color  and  odor  (?). 
Taste  a  drop  of  it  (  ?).  Rub  a  little  of  the  glycerine  between  the  fingers  ( ?). 
Test  it  with  litmus  paper  (?).  Ascertain  its  solubility  in  (a)  water, 
(b)  alcohol,  (c)  ether  and  (d)  chloroform  (?). 

Heat  a  little  glycerine  in  a  dry  test  tube  and  try  to  boil  it.  What  evi- 
dence have  you  that  it  undergoes  decomposition  ?  Can  glycerine  be 
distilled  by  itself  ?  How  may  it  be  distilled  (volatilized)  ?  Enumerate 
several  of  its  uses.  Write  the  rational  formula  for  glycerine. 

Glycerine  (glycerol),  C3H5(OH)3,  is  an  alcohol  containing  three  hydroxyl 
groups.  It  is  the  trihydric  or  triatomic  alcohol  of  the  radical  glycerl, 
C3H5,  formed  by  the  removal  of  three  atoms  of  hydrogen  from  the  satu 
rated  hydrocarbon,  propane  (C3H8),  and  combination  of  the  radical 
with  three  hydroxyl  groups.  The  common  animal  and  vegetable  fats 
and  oils  contain  glycerine  in  combination  with  the  fatty  acids.  These 
compounds  form  a  class  of  esters  known  as  glycerides,  which  when  treated 
with  alkalies  undergo  decomposition — the  fatty  acids  combining  with  the 
metals  of  the  alkali  to  form  soaps,  wrhilst  glycerine  is  liberated. 

Pure  glycerine  is  a  thick,  colorless,  odorless  liquid,  oily  to  the  touch, 
hygroscopic,  neutral  in  reaction  and  rather  swreet.  It  is  soluble  in  water 
and  alcohol  in  all  proportions,  but  insoluble  in  chloroform  and  ether. 
It  has  a  sp.  gr.  of  1.225,  and  solidifies  at  low  temperatures  forming  deli- 
quescent crystals  which  melt  at  17°  C.  When  heated  to  the  boiling-point, 
under  ordinary  atmospheric  pressure,  it  undergoes  decomposition;  there- 
fore, it  cannot  be  distilled  by  itself.  However,  it  can  be  distilled  under 
diminished  pressure  and  is  volatilized  in  the  presence  of  water  or  when 
hot  steam  is  allowed  to  pass  through  it. 

On  account  of  its  solvent  properties  glycerine  is  used  extensively  in  the 
preparation  of  official  solutions  of  various  organic  and  inorganic  sub- 
stances. These  solutions  have  been  termed  glycerites. 

"Nitro-glycerine"  C3H5(NO3)3  is  prepared  by  treating  glycerine  with 
a  mixture  of  concentrated  sulphuric  and  nitric  acids.  The  chemical 
action  is  indicated  by  the  following  equation: 

C3H5(OH)3  +  3HN03  -  C3H5(N03)3  +  3H2O  -  (  -  ?  cal.) 

It  is  a  pale-yellow  oily  liquid  soluble  in  alcohol  but  insoluble  in  water; 
crystallizes  at  — 20°  C.  in  long  needles  and  explodes  very  violently  by 
concussion.  It  can  be  burned  in  an  open  dish,  but  if  heated  above  250°  C. 
it  explodes.  One  kilogram  of  nitro-glycerine  yields  after  explosion  about 
725  liters  of  gas,  measured  at  o°  C.,  and  760  mm.  pressure.  As  the 
temperature  of  the  gas  is  raised  to  about  7000°  C.  by  heat  liberated  by 
the  explosion,  the  volume  is  much  larger  than  that  suggested.  The 
explosive  power  of  nitro-glycerine  is  about  thirteen  times  as  great  as  that 
of  gunpowder.  Nitro-glycerine  is  the  active  constituent  of  a  number 
of  explosives.  Dynamite  is  infusorial  earth  impregnated  with  nitro- 


256  EXPERIMENTAL  CHEMISTRY. 

glycerine.     While  it  is  not  readily  exploded  by  pressure  or  jar,  it  is  by 
percussion. 

Experiment  VI. — Properties  of  Phenol  (Carbolic  Acid).     Tests. 

Caution. — Phenol  is  strongly  poisonous,*  and  causes  blisters  if  it 
comes  in  contact  with  the  flesh. 

(a)  Ask  the  instructor  to  show  you  a  specimen  of  solid  phenol  or  a 
solution  of  it.     Observe  its  empyreumatic  and  disagreeable  odor.     When 
diluted  greatly,  it  possesses  a  sweetish  and  afterward  a  caustic  taste.     Test 
a  dilute  solution  with  litmus  paper  (?).     Recall  its  solubility  in  water 
(See  Exp.  XI,     "Solutions").     What  peculiar  property  of  solubility  in 
water  does  it  possess?     Define  "critical  solution  temperature."     What 
is  the   melting  point  of  phenol?     Indicate   by  rational  formula  that 
phenol  (phenyl  hydrate)  has  the  structure  of  an  alcohol. 

(b)  To  2  cm. ^  of  phenol  add  small  quantities  of  a  solution  of  potassium 
hydroxide  until  the  phenol  is  in  solution.     What  are  the  products  of  the 
interaction?     Now    add    an    excess    of    hydrochlorid    acid.     Results? 
Equations  ? 

(c)  Tests. 

1.  Add   a  few  drops  of   a  neutral  solution  of   ferric  chloride  to  an 
aqueous  solution  of  phenol.     A  beautiful  blue  color  is  imparted  to  the 
solution. 

2.  Bromine  water  added  in  excess  gives  a  yellowish-white  precipitate 
of  tri-brom-phenol,  C6H2Br3OH,  which  has  been  used  medicinally  under 
the  name  of  "bromol." 

3.  When  phenol  is  heated  with  nitric  acid  it  turns  yellow  owing,  to  the 
formation  of  picric  acid  (trinitro-phenol),  C6H2(NO2)3OH. 

Phenol  (carbolic  acid,  phenyl  hydrate,  phenyl  alcohol),  C6H5OH,  is 
prepared  by  the  distillation  of  coal  tar.  f  It  is  one  of  the  chief  constituents 
of  the  distillate  obtained  between  I7o0-i9o0  C.  When  the  distillate  is 
chilled  the  naphthalene  crystallizes  out,  leaving  the  phenol  mixed  with 
various  neutral  oils  and  impurities. 

The  remaining  liquor  is  treated  with  an  alkali  which  dissolves  the 
phenol,  forming  a  solution  of  "sodium  carbolate"  which  separates  by 
.gravity  from  the  undissolved  neutral  oils.  From  the  solution  it  is  pre- 
cipitated by  sulphuric  acid,  carbon  dioxide  or  furnace  gases — the  crude 
carbolic  acid  separating  as  an  oily  liquid.  It  is  purified  by  repeated 
distillations  and  crystallizations. 

Pure  carbolic  acid  is  obtained  in  the  form  of  colorless,  interlacing, 
needle-shaped  crystals  which  sometimes  acquire  a  pinkish  tint.  The 
presence  of  water  prevents  it  from  solidifying.  It  melts  at  43°  C.  and 
boils  at  183°  C.  The  sp.gr.  is  1.065.  It  is  very  soluble  in  alcohol, 

*  Antidote. — Castor  oil,  or  olive  oil,  or  a  mixture  of  both,  or  a  mixture  of  magnesia 
and  oil,  also  sodium  sulphate  (Glauber's  salt),  internally  and  hypodermically  are  used 
as  antidotes. 

t  Thorpe's  Outlines  of  Industrial  Chemistry. 


SOME    COMMON   CARBON   COMPOUNDS.  257 

ether  and  glycerine.  At  17°  C.,  it  dissolves  in  20  parts  of  water.  It 
also  dissolves  a  small  amount  of  water;  therefore  if  phenol  is  shaken  with 
less  than  20  parts  of  water  at  17°  C.,  the  mixture  will  separate  into  an 
upper  layer  containing  about  five  per  cent,  and  a  lower  layer  containing 
about  seventy-five  per  cent,  of  phenol.  At  or  above  68.9°  C.,  the  two  are 
miscible  in  all  proportions.  (Recall  Exp.  XI,  "Solutions.")  It  is 
very  poisonous  and  a  very  powerful  disinfectant  or  germicide. 

Phenols  are  the  hydroxyl  derivatives  of  benzene,  C0H6.  Their  con- 
stitution allies  them  with  the  alcohols,  but  they  possess  properties  so 
different  from  those  of  the  alcohols  of  the  methane  and  ethylene  series 
that  they  have  been  given  the  general  but  distinctive  name  by  which  they 
are  now  known.  They  differ  from  the  common  alcohols  in  not  yielding 
aldehydes  or  acids  by  oxidation.  Although  phenol  is  generally  called 
carbolic  acid,  it  has  a  neutral  or  but  faintly  acid  reaction.  It  combines 
with  strong  bases  to  form  a  class  of  salts  known  as  the  carbolates  or 
phenolates. 

C0H5OH  +  NaOH ->  C6H5ONa  +  H2O. 
Phenol  Sodium 

Carbolate 

C6H5OH  +  HC?H5ps—  C0H5.C7H5O3  +  H2O. 
Phenol  Salicylic         Salol 

Acid 

2C6H00  +  C8H403->  C20H1404  +  H20. 
Phenol          Phthalic      Phenol- 
anhydride         phthalein. 

Experiment  VII. — Preparation  and  Properties  of  Ether. 

Caution. — In  working  with  ether  carefully  avoid  the  neighborhood 
of  free  flames  as  its  vapor  is  easily  ignited. 

(a)  Preparation.     Recall  or  repeat  Exp.  IV  (c).     Equation? 

(b)  Pour  a  few  cm.3   of    ether  into  a  test  tube.     Observe  its  color, 
odor  and  volatility  (?).     Is  it  lighter  or  heavier  than  water?     Taste  a 
drop   of    it  (?).     Allow    a  few    drops    to    fall    upon    the    back    of    the 
hand    (?).      Pour    a   drop   upon   a   watch   glass.     Does    it    evaporate 
more  or  less  rapidly  than  alcohol?     Predict  its  boiling-point  relative  to 
that  of  alcohol  (?).     Reasons  for  your  answer?     Test  the  ether  with 
litmus  paper   (?).     What  can  you   say  with  reference   to   the   mutual 
solubility  of  ether  and  water?     (Recall  Exp.  X,  "Solutions.") 

To  a  small  piece  of  resin  or  wax  add  several  cm.3  of  ether  (?).  The 
result  is  typical.  What  are  your  conclusions  as  to  its  solvent  properties  ? 
Write  the  rational  formula  for  ether  (ethyl).  To  what  class  of  inor- 
ganic compounds  is  it  analogous? 

Our  ordinary  ether  (ethyl  ether)  is  the  representative  of   a  very  large 
class  of  organic  compounds  which  are  given  the  general  name  of  ethers. 
These  compounds  are  the  oxides  of  the  organic  radicals  and  are  analogous 
17 


258  EXPERIMENTAL  CHEMISTRY. 

to  the  various  metallic  oxides.  Many  ethers  are  products  of  vegetable 
life,  and  have  generally  a  characteristic  and  pleasant  odor.  Fruit 
essences  consist  mainly  of  compound  ethers  (esters).  That  which  is 
known  as  the  "bouquet"  or  " flavor"  of  wine  and  other  alcoholic  liquors 
is  due  chiefly  to  ethers  or  compound  ethers,  which  are  formed  during  the 
various  stages  of  fermentation  by  the  action  of  the  acids  present  upon  the 
alcohol  or  alcohols  formed.  The  improvement  which  such  alcoholic 
liquids  undergo  "by  age"  is  caused  by  a  continued  chemical  action  be- 
tween the  substances  named.  Pure  ethyl  ether  is  a  colorless,  mobile 
liquid  which  has  a  sp.gr.  of  0.718  at  156°  C.  It  boils  at  34.6°.  On 
account  of  its  volatility,  ether  should  be  kept  in  strong  bottles  tightly 
corked.  It  is  easily  combustible  and  burns  with  a  luminous  flame. 
Ether  mixes  with  alcohol  in  all  proportions.  Official  ether  contains  96 
per  cent,  of  ethyl  ether  and  4  per  cent,  of  alcohol.  It  is  stated  by  Noyes 
(W.  A.)  that  ether  "dissolves  in  n.i  volumes  of  water  at  25°  C.,  while  it 
will  in  turn  dissolve  one-fiftieth  of  its  volume  of  \vater."  It  is  a  good 
solvent  for  fats,  alkaloids,  resins  and  many  other  classes  of  organic  com- 
pounds. It  is  used  extensively  as  an  anesthetic,  causing  intoxication 
and  finally  loss  of  consciousness  and  sensation  when  inhaled.  It  is 
neutral  in  reaction. 

Experiment  VIII. — Oxidation  of  an  Alcohol  to  an  Aldehyde. 

(a)  Acetic  or  ethyl  aldehyde.     Recall  or  repeat  Exp.  IV  (b).     Equa- 
tion? 

C2H6O  — 2H  -»  C2H4O 
Ethyl  Acetic 

alcohol  aldehyde. 

(b)  Formic  aldehyde  or  formaldehyde.     This  substance  is  produced 
by  passing  a  mixture  of  the  vapor  of  methyl-alcohol  and  air  over  a  heated 
copper  spiral.     The  experiment  may  be  performed  as  follows:     Pour 
5  cm.3  of  methyl-alcohol  into  a  test  tube  clamped  in  a  vertical  positon — 
then  drop  into  the  alcohol  a  spiral  of  copper  wire  which  has  been  heated 
to  redness  in  a  flame.     The  colorless  penetrating  gas  wrhich  is  evolved 
is  formaldehyde.     What  is  the  formula  for  formaldehyde  ? 

CH3OH  +  O  -*  H,CHO  +  H2O. 

Aldehydes  are  formed  by  the  removal  of  hydrogen  from  alcohols. 
In  fact,  the  name  aldehyde  is  derived  from  alcohol  dehydrogenatum  which 
refers  to  its  method  of  formation.  The  removal  of  hydrogen  may  be 
accomplished  by  various  methods,  as,  for  example,  by  the  oxidation  of 
alcohols,  when  one  atom  of  oxygen  combines  with  two  atoms  of  hydrogen, 
forming  water,  while  an  aldehyde  is  formed  simultaneously.  The 
realtions  in  composition  of  the  alcohols,  aldehydes  and  acids  are  shown 
by  the  following  formulae: 

Radicals      Hydrocarbons     Alcohols      Ethers     Aldehydes     Acids 
CH3  CH4  CH3OH    (CH3)2O       C  H2O       C  H2O2 

C2H5  C2H6  C2H5OH    (C2H5)20       C2H4O       C2H4O2 


SOME  COMMON  CARBON  COMPOUNDS.  259 

Aldehydes,  when  further  oxidized,  are  converted  into  acids.  Many 
of  the  aldehydes  in  consequence  of  their  tendency  to  unite  with  oxygen 
to  form  acids  are  strong  reducing  agents.  Only  a  few  of  the  aldehydes 
are  of  practical  interest,  as,  for  examples,  formaldehyde  (CH2O  or 
H.COH),  acetaldehyde  (C2H4O  or  CH8.COH),  the  polymerized  form 
of  acetaldehyde,  paraldehyde  ((C2H4O)3)  and  a  few  others. 

Formaldehyde  is  a  gas  at  ordinary  temperatures  but  may  be  condensed 
to  a  liquid  which  boils  at  — 21°  C.  It  is  readily  soluble  in  water,  and  a 
40  per  cent  solution  has  been  placed  on  the  market  under  the  name  of 
jormalin.  This  same  name  is  now  given  to  solid  paraformaldehyde, 
(CH2O)3,  which  is  a  polymerized  form  of  formaldehyde.  It  is  extensively 
used  as  an  antiseptic.  The  paraformaldehyde  (formalin)  splits  up  into 
three  molecules  of  formaldehyde,  which,  escaping  as  a  gas,  is  used  for 
disinfecting  purposes.  It  does  not  act  injuriously  on  the  fabric  or  color 
of  household  goods,  thus  possessing  an  advantage  over  chlorine  and 
sulphur  dioxide. 

Experiment  IX. — Organic  Acids.     Acetic  Acid. 

(a)  To  a  solution  of  sodium  acetate,  NaC2H3O2,  add  a  little  dilute  sul- 
phuric acid.     Observe  the  odor — it  may  be  necessary  to  gently  warm  the 
mixture  (?).     Equation?     Name  two  commercial  processes  by  which 
acetic  acid  is  manufactured. 

(b)  Test  a  solution  of  acetic  acid  with  litmus  paper  ( ?).     Pour  a  few 
drops  of  the  solution  upon  a  solution  of  sodium  carbonate,  Na^COg. 
Results  ?     Equation  ?     What  per  cent,  is  ionized  in  a  i  N  solution  ?     In  a 
.1  N  solution?     (See  table).     Would  you  infer  that  it  is  a  "strong"  or 
a  "weak"  acid?     Test  its  relative  electrical  conductivity  (?).     What 
class  of  salts  are  formed  by  acetic  acid  ? 

(c)  Compare  the  odor  of  dilute  acetic  acid  with  vinegar  (?).     Try 
the  action  of  vinegar  upon  a  solution  of  sodium  carbonate  ( ?).     Equation  ? 
What  are  your  conclusions  as  to  the  relation  of  acetic  acid  to  vinegar? 

(d)  Tests  for  acetic  acid. 

1.  Acetates  heated  with  sulphuric  acid  evolve  acetic  acid  which  may 
be  readily  detected  by  its  odor. 

2.  Acetates   or   acetic  acid   heated  with  sulphuric  acid  and  alcohol 
yield  the  characteristic  odor  of  ethyl  acetate.     (See  Exp.  IV.  (d)). 

3.  Ferric  chloride,  when  added  to   a  solution  containing  acetic  acid 
or  an  acetate  carefully  neutralized,  gives  ferric  acetate,  which  is  a  soluble 
salt  and  imparts  a  dark  red  color  to  the  solution.     When  this  solution  is 
boiled,  a  brownish  precipitate  of  basic  ferric  acetate  separates  out. 

As  suggested  previously,  many  organic  acids  are  produced  by  the  oxi- 
dation of  alcohols.  They  show  the  characteristics  mentioned  of  inorganic 
acids,  viz.,  when  soluble,  have  an  acid  or  sour  taste,  redden  litmus,  con- 
tain hydrogen  which  is  replacable  by  metals  with  the  formation  of  salts, 
and  yield  hydrogen  ions.  The  greater  number  of  these  organic  acids 
possess  these  acid  properties  in  a  much  less  marked  degree;  in  fact,  they  are 


260  EXPERIMENTAL  CHEMISTRY. 

so  weak  that  the  acid  properties  can  often  be  scarcely  detected.  Most 
organic  acids  are  colorless,  odorless  solids;  a  few  are  liquids  and  scarcely 
any  are  gaseous  at  the  ordinary  temperature. 

Pure  acetic  acid,  or  glacial  acetic  acid,  melts  at  16.7°  C.  and  boils  at 
120°  C.  It  has  a  sp.  gr.  of  1.055.  It  causes  blisters  on  the  skin  and  is 
miscible  in  water,  alcohol  and  ether.  It  forms  salts  known  as  acetates, 
all  of  which  are  soluble  in  water.  (See  Thorpe's  "  Outlines  of  Industrial 
Chemistry"  and  Sadtler's  "Industrial  Organic  Chemistry"  for  the 
manufacture  of  acetic  acid.) 

Experiment  X. — Tartaric  Acid.     Tests. 

(a)  Examine     some    crystals    of     tartaric     acid.      Is    it    soluble    in 
water?     What  is  the  effect  of  the  solution  on  blue  litmus  paper?     Try 
the  effect  of  an  aqueous  solution  of  tartaric  acid  upon  a  solution  sodium 
bicarbonate  ( ?).     Equation  ? 

(b)  Give  the  formula  and  state  one  well-defined  use  of  each  of  the 
following:     cream  of  tartar,  Rochelle  salt,  tartar  emetic. 

(c)  Tests. 

1.  Tartrates  are  readily  charred  by  heating  them  with  strong  sulphuric 
acid. 

2.  Tartrates  are  decomposed   (char),  and  evolve  an  odor  resembling 
that  of  burnt  sugar  when  heated. 

3.  Silver   nitrate   gives  with   a  neutral  solution  of  a  tartrate  a  white 
precipitate  of  silver  tartrate  which  blackens  on  boiling  in  consequence  of 
the  decomposition  of  the  salt,  with  separation  of  silver.     If  ammonium 
hydroxide  is  added  before  boiling,  a  mirror  of  metallic  silver  will  form 
upon  the  glass. 

Tartaric  acid,  H2C4H4O6,  is  frequently  found  in  vegetables  and  fruits, 
and  especially  in  grapes.  With  the  exception  of  the  tartrates  of  the  alka- 
lies, all  normal  tartrates  are  either  insoluble  or  difficultly  soluble.  Most 
of  the  insoluble  tartrates  dissolve  in  tartaric  acid  forming  "acid"  salts. 

Experiment  XI.— Oxalic  Acid.     Tests. 

Caution. — Oxalic  acid  is  a  poison. 

(a)  Repeat  Exp.  X  (a),  substituting  oxalic  acid  for  tartaric  acid. 

(b)  Place  5  grm.  of  sugar  in  150  cm.3  flask;  add  50  cm.3  of  nitric  acid 
(i  -  i);  heat  gently  in  the  hood  until  brown  fumes  are  given  off,  then  re- 
move the  flame.     When  the  brown  fumes  cease  to  escape,  boil  the  liquid 
down  rapidly  to  about  10  or  15  cm.3;  pour  the  solution  into  a  crystallizing 
dish.     Crystals  of  oxalic  acid  separate  on  cooling.     Remove  the  crystals 
and  dissolve  them  in  warm  water  and  recrystallize  them.      Dissolve  some 
of  these  crystals  in  water,  and  add  this  solution  to  a  few  cm.3  of  a  dilute 
solution  of  potassium  permanganate  which  has  been  acidulated  with  a  few 
drops  of  sulphuric  acid.     Is  the  permanganate  decolorized  ?     Equation  ? 

(c)  Tests. 

i.  When  strong  sulphuric  acid  is  heated  with  oxalic  acid  or  oxalates 
carbon  monoxide  and  carbon  dioxide  are  evolved. 


SOME  COMMON  CARBON  COMPOUNDS.  261 

2.  Calcium  chloride  gives  with  neutral  solutions  of  oxalic  acid  a  white 
precipitate  of  calcium  oxalate,  CaC2O4,  which  is  soluble  in  hydrochloric, 
but  insoluble  in  acetic  acid. 

Oxalic  acid,  H2C2O4,  is  a  white  solid,  soluble  in  water,  crystallizing 
from  the  solution  with  two  molecules  of  water  of  crystallization,  H2C2O4, 
2H2O.  It  occurs  very  widely  distributed  in  the  vegetable  kingdom,  as 
in  certain  plants  of  the  oxalis  varieties,  in  the  form  of  the  acid  potassium 
salt.  It  may  be  obtained  by  the  oxidation  of  many  organic  substances, 
chiefly  sugars,  starches,  etc.,  by  nitric  acid  or  other  strong  oxidizers.  It 
is  prepared  technically  by  heating  sawdust  or  wood-shavings  with  a 
mixture  of  caustic  soda  and  caustic  potash  to  about  250°  C.,  when  the 
oxalate  of  these  metals  are  formed.  The  mass  is  extracted  with  water, 
and  the  solution  evaporated  to  crystallization,  when  sodium  oxalate  is 
deposited  or  the  dissolved  alkali  oxalate  is  treated  with  calcium  hydrox- 
ide, insoluble  calcium  oxalate  forming.  The  calcium  oxalate  is  de- 
composed by  sulphuric  acid. 

Oxalic  acid  is  probably  the  strongest  organic  acid.  It  is  dibasic  and 
acts  as  a  reducing  agent,  decolorizing  solutions  of  the  permanganates. 
It  possesses  poisonous  properties.* 

Oxalic  acid  is  largely  used  in  dyeing,  calico-printing  and  bleaching, 
in  cleaning  brass  and  in  removing  iron-mould  from  linen. 

The  oxalates  of  the  alkalies  are  soluble  in  water.  All  other  oxalates  are 
insoluble  or  only  sparingly  soluble. 

Experiment  XII. — Halogen  Derivatives.     lodoform.     Chloroform. 

(a)  lodoform.     Recall  or  repeat  Exp.  IV  (a).     Equation? 

(b)  Chloroform.      The  following  method  for  its  preparation  is  sug- 
gested by  Remsen:      Mix  550  grams  of  bleaching  powder  and  1.25  liters 
water  in  a  3-liter  flask.     Add  33  grams  alcohol  of  sp.  gr.  0.834.     Heat 
gently  on  a  water  bath  until  action  begins.     A  mixture  of  alcohol,  water 
and  chloroform  will  distill  over.     Add  water  and  remove  the  chloroform 
by  means  of  a  pipette.     Add  calcium  chloride  to  the  chloroform  and, 
after  standing,  distill  on  a  water  bath. 

4C2H6O  +  8  Ca(ClO)2—  2CHC13  +  3(Ca(CHO2)2)  +  5CaCl2  +  8  H2O. 

(c)  Note  the  odor  of  pure  chloroform.     Taste  a  drop  of  it  (?).     Test 
its  reaction  with  litmus  paper  (?).     Shake  a  little  chloroform  with  a 
few  cm.3  of  a  solution  of  silver  nitrate  (?).     Name  one  of  the  principal 
uses  of  chloroform. 

(d)  Tests  for  chloroform. 

1.  Add  a  few  drops  of  chloroform  to  5  cm.3  of  Fehling's  solution  and 
heat;  red  cuprous  oxide  is  thrown  down. 

2.  A  strip    of    filter   paper  saturated  with  chloroform  burns  with  a 
green-mantled    flame   which    emits  vapors  of  hydrochloric   acid.     The 

*  Antidote. — "Lime-water  or  calcium  carbonate  should  be  administered,  but  no 
alkalies  as  in  cases  of  poisoning  by  mineral  acids,  because  alkali  oxalates  are  soluble." 
— Simon's  Manual  of  Chemistry. 


262  EXPERIMENTAL  CHEMISTRY. 

acid  vapors  may  be  rendered  more  visible  by  approaching  the  flame 
with  a  glass  rod  moistened  with  ammonium  hydroxide. 

Chloroform  is  a  heavy,  colorless  liquid,  possessing  a  sp.  gr.  of  1.526 
at  o°  C.;  it  has  a  burning  sweet  taste;  a  neutral  reaction,  a  boiling-point 
of  61.2°  C.,  and  a  melting-point  of  — 70°  C.  It  is  very  slightly  soluble  in 
water,  but  miscible  with  alcohol  and  ether  in  all  proportions.  Pure  chlo- 
roform is  not  very  stable.  It  undergoes  decomposition  slowly  when  ex- 
posed to  light  and  air,  yielding  hydrochloric  acid,  free  chlorine,  phosgene 
(COC12)  and  other  products.  Some  of  the  decomposition  products  are  more 
poisonous  than  chloroform,  therefore  when  the  latter  is  used  as  an  anes- 
thetic it  is  essential  that  the  chloroform  should  be  pure.  It  has  been 
found  that  the  presence  of  a  small  amount  (.5  to  i%)  of  alcohol  renders 
the  chloroform  more  stable.  Commercial  chloroform  usually  contains 
a  small  quantity  (the  U.  S.  P.  allows  one-half  to  one  per  cent.)  of  al- 
cohol for  this  purpose;  thereby  changing  the  sp.  gr.  to  about  1.488. 
Pure  chloroform  should  not  give  more  than  a  faint  opalescence  when 
shaken  with  a  silver  nitrate  solution. 


Experiment  XIII. — Esters.     Ethyl  Acetate. 

Recall  or  repeat  Exp.  IV  (d).     Equation? 


Recall  the  note  following  Exp.  IV.  The  esters  are  analogous  to  inor- 
ganic salts  in  structure,  and  it  is  frequently  convenient  to  refer  to  them  as 
"organic  salts,"  but  their  resemblance  to  said  salts  ceases,  however,  with 
structure.  The  esters  react  very  slowly.  Originally,  the  esters  were 
called  "compound  ethers";  for  example,  ethyl  acetate  was  called  "acetic 
ether."  The  terms  are  still  occasionally  employed. 
,;  All  true  jats  are  esters  of  the  triatomic  alcohol,  glycerine,  and  the  higher 
members  of  the  fatty  acids.  They  are  frequently  termed  glycerides  ;  stearin, 
for  example,  being  the  glyceride  of  stearic  acid.  Fats  are  widely  dis- 
tributed in  the  animal  and  vegetable  kingdoms.  They  are  found  in 
animals  generally  under  the  skin,  around  the  intestines  and  on  the  muscles, 
while  in  plants  they  exist  chiefly  in  the  seeds. 

Human  fat,  lard,  mutton  fat  and  bee/  tallow  are  mixtures  of  palmitin, 
C3H5(C16H3102)3,  stearin,  C3H5(C18H35O2)3,  and  olein,  C3H5(C18H33O2)3. 

Butter  consists  essentially  of  the  glycerides  of  butyric  acid,  caprylic  acid, 
caproic  acid  and  capric  acid  which  are  volatile  with  steam,  and  of  palmitic 
and  stearic  acids,  which  are  not  volatile.  Palmitin  and  stearin  are  solids, 
but  olein  is  a  liquid  at  ordinary  temperatures.  The  relative  amount  of 
these  three  fats  determines  its  solid  or  liquid  condition.  The  chief  con- 
stituent of  liquid  fats  is  generally  olein.  The  liquid  fats  have  been  given 
the  name  jatty  oils  or  fixed  oils  in  contradistinction  to  the  essential  or 
volatile  oils  which  give  the  peculiar  odors  to  plants.  Among  the  more 
familiar  essential  oils  are  spirits  of  turpentine,  camphor,  menthol,  cedar 
oil,  oils  oj  cloves,  lemon,  mustard,  peppermint,  and  wintergreen,  attar  oj 
roses,  etc. 


SOME  COMMON  CARBON  COMPOUNDS.  263 

The  taste  and  color  of  fats  are  due  to  foreign  substances  which  are 
often  produced  by  slight  decomposition.  Pure  fats  are  odorless,  colorless 
and  tasteless  substances  which  stain  paper  permanently.  They  are  lighter 
than  water  and  insoluble  in  it,  sparingly  soluble  in  alcohol,  easily  soluble 
in  ether,  carbon  disulphide,  benzene,  etc.  Fats  heated  above  300°  C. 
decompose  with  the  formation  of  various  products,  some  of  which  possess 
a  very  disagreeable  odor.  Among  the  products  is  the  aldehyde,  acrolein, 
C,H40. 

Many  of  the  pure  fats  keep  without  change,  but  since  the  greater  num- 
ber contain  impurities,  such  as  albuminous  matter,  etc.,  they  suffer  de- 
composition,* which  results  in  a  liberation  of  the  fatty  acids.  The  latter 
impart  their  odor  and  taste  to  the  fats,  causing  them  to  become  what  is 
generally  termed  rancid. 

Drying  oils  are  a  class  of  fats  containing  unsaturated  acids  which 
undergoing  oxidation,  render  them  hard.  The  principal  vegetable  dry- 
ing oils  are  linseed  oil,  hemp  oil,  poppy  oil  and  sunflower  oil. 

Among  the  vegetable  semi-drying  oils  are  corn  or  maize  oil,  cotton-seed 
oil  and  castor  oil. 

The  vegetable  non-drying  oils  of  importance  are  olive  oil,  peanut  oil, 
palm  oil,  cocoanut  oil  and  almond  oil. 

Butterine  and  oleomargarine  are  perfect  substitutes  for  common  butter 
when  the  essential  attributes  are  considered.  They  are  made  from  mix- 
tures of  animal  and  vegetable  oils  flavored  with  a  relatively  small  quantity 
of  butter  and  colored  to  imitate  it.  Oleo  oil  from  tallow  and  lard  are 
largely  used.  These  are  frequently  mixed  with  cotton-seed  oil  in  cold 
weather  to  increase  the  percentage  of  olein. 

Experiment  XIV. — Preparation  of  Soaps.     Glycerine. 

(a)  Preparation  of  soda  soap.     Put  about  10  grams  of  lard  or  tallow 
in  a  300  cm.3  Erlenmeyer  flask;  dissolve  4  grams  of  NaOH  in  100  cm.3 
of  water,  pour  this  solution  into  the  flask;  heat  the  mixture  over  a  wire 
gauze  until  it  ceases  to  foam,  then  add  an  equal  volume  of  a  saturated 
solution  of  sodium  chloride.     Stir  constantly  during  the  addition  of  the 
salt  solution.     Boil  for  a  few  minutes.     Allow  the  mixture  to  stand  until 
it  separates  into  an  upper  solid  zone  or  crust  and  a  lower  liquid  zone 
which  may  be  more  or  less  colored.     The  solid  layer  is  a  mixture  of  soda- 
soap  and  unsaponified  fat;  the  liquid  contains  glycerine,  salt  and  various 
impurities.     Drain  off  this  liquid  and  add  3  grams  of  NaOH  dissolved  in 
75  cm.3  of  water;  heat  as  previously  directed  until  the  mixture  becomes 
adhesive.     Pour  the   mixture  into   an   open  dish.     The   solid  product 
should  show  the  essential  properties  of  hard  soap.     Equations?     Explain 
the  precipitation  ("salting  out")  of  the  soap  by  means  of  sodium  chloride. 

(b)  Preparation  of  potassium  soap.     Dissolve  5  grams  of  KOH  in  100 
cm.3  of  alcohol;  add  7  grams  of  lard;  heat  the  mixture  on  a  water  bath  until 
it  has  the  consistency  of  syrup  and  the  odor  of  alcohol  is  no  longer  per- 
ceptible.    The  mixture  should  be  stirred  constantly  while  it  is  being 

*  A  kind  of  fermentation  which  is  aided  by  oxidation. 


264 


EXPERIMENTAL  CHEMISTRY. 


heated.  Pour  the  mixture  into  an  open  dish  and  allow  it  to  cool.  The 
jelly-like  product  is  soft  soap.  The  alcohol  is  used  as  a  common  solvent 
for  the  fat  and  alkali. 


C3H5(C16H31O2)3 
Palmitin 
(Fat) 

C3H5(C18H35p2)3 
Stearin 
(Fat) 

C3H5(C18H33O2)3 
Olein 

C3H5(C16H31O2)3 

Palmitin 


3NaOH   -*  3NaC16H3102. 

Sodium  palmitate 
(Hard  soap) 


+  3NaOH   -.    3NaC18H3602 

Sodium  stearate 
(Hard  soap) 

+  3NaOH   -»  3NaC18H33O2 
Sodium  oleate 

+  3KOH     ^3KC16H3102 

Potassium  palmitate 
(Soft  soap) 


C3H5(OH)3 
Glycerin 

C3H5(OH)3 
Glycerin 

C3H5(OH)3 
Glycerin 

C3H5(OH)3 
Glycerin 


Experiment  XV.  —  Properties  of  Soaps. 


(a)  Dissolve  a  little  of  the  soap  prepared  in  Exp.  XIV  (a)  in  warm 
water;  filter.     To  one-half  of  the  filtrate  add  hydrochloric  acid  and  shake 
vigorously.     Observe  that  the  fatty  acids  separate  as  solids  and  rise  to 
the  top.     Remove    this    floating    coagulum    and    test    its  solubility  in 
alkali  (?).     Equations? 

(b)  To  the  other  half  of  the  soap  solution  add  a  calcium  sulphate  or 
magnesium  sulphate  solution.     Note  the  formation  of  the  corresponding 
insoluble   calcium   or   magnesium   salt   of   the   fatty  acid.     Equations? 
Explain  how  the  efficiency  of  soap  is  diminished  by  using  it  with  hard 
water.     What  information  do  we  convey  by  the  expression  sojt  water  ? 

(c)  Test  yellow  soaps,  toilet  soaps,  etc.,  for  free  alkali.     Tabulate  your 
results  giving  the  name  of  the  soap  used. 

As  previously  explained,  the  common  fats,  glycerides  or  esters  undergo 
decompostion  when  boiled  with  water,  acids  or  alkalies  writh  the  forma- 
tion of  the  free  acid  or  a  salt  of  the  acid,  and  the  alcohol  (glycerin). 
This  process  of  decomposing  an  ester  is  called  saponification.  The 
metallic  salts  which  are  formed  when  the  saponification  is  effected  by 
means  of  an  alkali  are  called  soaps.  They  are  defined  as  metallic  salts  of 
certain  non-volatile  fatty  acids.  Soaps  intended  for  washing  purposes 
should  contain  only  soluble  salts  of  the  acids;  i.e.,  those  of  sodium  or 
potassium;  the  magnesium,  calcium,  lead  and  other  heavy  metal  salts 
are  insoluble  in  water.  The  alkalies  which  are  most  frequently  used  in 
the  manufacture  of  soaps  are  sodium  and  potassium  hydroxides.  The 
former  yields  a  "hard  soap";  the  latter,  a  "soft  soap"  which  is  liquid 
under  ordinary  conditions  because  of  a  greater  solubility  and  a  lower 
melting-point.  It  is  also  probable  that  the  greater  deliquescence  of  the 
potassium  soaps  contributes  to  their  "softness." 


SOME  COMMON  CARBON  COMPOUNDS.  265 

Experiment  XVI. — Carbohydrates.     Grape  Sugar  (Glucose). 

(a)  Examine  some  glucose  (grape  sugar)  and  tabulate  its  properties. 
What  is  its  formula? 

(b)  Glucose   ferments.     Recall   or   repeat    Exp.    Ill  (a).   Equation? 

(c)  "Fehling's  Test"  for  glucose.     Place  about  10  cm.3  of  Fehling's 
solution  in  a  test  tube;  add  a  few  drops  of  a  solution  of  glucose;  boil  hard 
for  several  minutes,  then  add  a  little  more  of  the  glucose  solution.  Continue 
to  boil  until  on  removing  the  tube  from  the  flame,  a  dark-red  precipitate 
settles  to  the  bottom,  leaving  a  clear  colorless  liquid.     The  red  precipitate 
is  cuprous  oxide.      This  reaction  can  be  used  both  for  detection  and 
quantitative  determination,  if  other  reducing  substances  are  absent. 

Experiment  XVII. — Carbohydrates.  Cane  sugar  (Saccharose).  Inver- 
sion of  Cane  Sugar. 

(a)  Repeat  Exp.  XVI  (a),  substituting  cane  sugar  or  beet  sugar  for  glu- 
cose. 

(b)  Cane  sugar  does  not  ferment.     Repeat  Exp.  XVI  (b)  substituting 
cane  sugar  or  beet  sugar  for  glucose.     Results? 

(c)  Inversion  of  cane  sugar.     Prepare  a  dilute  solution  of  pure  cane- 
sugar  (rock-candy).     Divide  it  into  two  parts.     Test  one  portion  with 
Fehling's  solution  as  in  Exp.  XVI.     Result?     Now  add  a  few  drops 
of  concentrated  hydrochloric  acid  to  the  other  portion  of  the  sugar  solu- 
tion ;  heat  for  thirty  minutes  or  longer,  on  a  water  bath  at  i  oo  °  C . ;  make  the 
solution  exactly  neutral  with  a  solution  of  sodium  carbonate  and  test  with 
Fehling's   solution.     Results?     Conclusion   relative   to   the   presence   of 
glucose  ?     Equation  ?     What  is  the  function  of  the  hydrochloric  acid  ? 

Experiment  XVIII. — Carbohydrates.  Preparation  and  Properties  of 
Cellulose.  Parchment  Paper. 

(a)  Treat  a  small  quantity  of  cotton-wool  or  Swedish  filter  paper 
successively  with  ether,  alcohol,  water,  a  caustic  alkali,  and  finally,  a  di- 
lute acid.     Wash  with  water  (?).    Enumerate  some  of  the  properties  and 
uses  of   cellulose.     What  is  its  formula? 

(b)  Parchment  paper.     Immerse  a  strip  of  filter  paper  in  a  solution 
prepared  by  adding  3  cm.3  of  concentrated  sulphuric  acid  to  15  cm.3 
of  water.     Allow  the  paper  to  remain  in  the  solution  for  about  fifteen 
seconds,  then  remove  it  to  a  large  vessel  of  pure  water  and  wash  it  thor- 
oughly.     Rinse  it  in  a  very  dilute  solution  of  ammonia  and  dry  it  in  the 
air.     Result  ? 

Experiment  XIX. — Conversion  of  Cellulose  into  Glucose. 

Cut  a  filter  paper  into  small  pieces  and  dissolve  it  in  the  least  volume 
of  concentrated  sulphuric  acid;  add  500  cm.3  of  water  and  boil  for  about 
an  hour.  Add  pulverized  chalk  (calcium  carbonate)  to  remove  the  sul- 
phuric acid;  filter,  and  evaporate  the  filtrate  to  small  bulk.  Test  for 


266  EXPERIMENTAL  CHEMISTRY 

glucose  with  Fehling's  solution.     Result  ?     The  chemical  reaction  may  be 
indicated  by  the  following  equation: 

(C6H1005)n  +  nH20  -*  n(C6H1206) 
Cellulose  Glucose 

Experiment  XX.  —  L.  T.  Cellulose  Nitrate  (Gun-cotton  or  Smokeless 
Powder). 

(a)  Mix  i  volume  of  nitric  acid  which  has  a  sp.  gr.  of  1.4  with  three 
volumes  of  concentrated  sulphuric  acid;  cool  the  mixture  to  about  12°  C.; 
add  as  much  pure  cotton,  small  quantities  at  a  time,  as  can  be  sub- 
merged in  the  acid  mixture.  Let  the  cotton  remain  in  the  mixture  for 
twenty-four  hours,  then  remove  it  and  wash  it  several  times  in  large 
volumes  of  water.  Let  it  dry  in  the  air.  The  white,  friable  mass  which 
is  very  unstable  is  gun-cotton.  To  a  small  quantity  of  the  prepared 
explosive  touch  a  lighted  match  (  ?). 


2(C6H1005)  +  6HN03-^  C12H1404(N03)6  +  6H2O* 
Decomposition  on  explosion: 
C12H1404(N03)6—  7C02  +  6CO  +  5~N~+  8H~+  3H2O  +  1074  calf 

(b)  State  briefly  the  composition,  preparation  and  uses  of  each  of  the 
following:  cordite,  celluloid  and  collodion. 

Experiment  XXI.  —  Carbohydrates.     Properties  of  Starch. 

(a)  Dust  a  glass  slide  with  starch,  then  examine  the  starch  through  a 
microscope.     Make  a  diagram  of  the  "field  of  vision." 

(b)  Ascertain  if  starch  is  appreciably  soluble  in  cold  water  or  alco- 
hol (?).     In  hot  water  (?). 

(c)  Test  for  starch.     Recall  test  for  iodine  (?). 

(d)  Cut  a  very  thin  section  from  a  piece  of  potato  by  use  of  a  micro- 
tome or  a  razor;  float  the  section  in  a  small  dish  of  water  to  which  has  been 
added  a  few  drops  of  an  iodine  solution.     Remove  the  section  from  the 
water  and  spread  it  upon  a  glass  plate,  then  cover  it  with  a  "cover  glass." 
Examine  the  section  through  a  microscope.     Draw  a  diagram  of  what 
you  see. 

Experiment  XXII.  —  Conversion  of  Starch  into  Glucose. 

(a)  Test  a  little  starch  paste  with  Fehling's  solution  (?). 

(b)  To  100  cm.3  of  starch  paste  add  5  cm.3  of  concentrated  sulphuric 
acid;  heat  the  mixture  on  a  steam  bath  for  two  hours  or  until  iodine  no 
longer  causes  a  blue  color  (which  shows  complete  conversion  of  starch 
into  either  dextrin  or  glucose),  and  until  i  cm.3  of  the  solution  yields  no 
precipitate  when  5  cm.3  of  alcohol  are  added  to  it,  dextrin  being  pre- 
cipitated by  alcohol.     The   dextrin  should   be   converted  into  glucose. 

*  Guttmann,  Industrie  der  Explosivstoffe.  f  Berthelot. 


SOME   COMMON   CARBON   COMPOUNDS.  267 

If  the  flask  containing  the  mixture  is  connected  with  an  inverted  con- 
denser, it  will  avoid  the  necessity  of  replacing  the  water  which  would  other- 
wise escape.  Proceed  from  this  point  as  in  Exp.  XIX.  Results  ? 

(C6H1005)n  +  nH20  —  n(C6H1206). 
Starch  Glucose. 

Carbohydrates. — These  compounds  are  found  widely  disseminated  in 
the  vegetable  kingdom;  in  fact,  no  other  organic  substances  are  found  in 
such  abundance.  They  are  also  found  as  products  of  animal  life,  as, 
for  example,  the  sugar  in  milk. 

Most  carbohydrates  are  white,  solid  substances,  generally  soluble 
in  water.  Those  belonging  to  the  sugars  have  a  more  or  less  sweet  taste. 
They  are  of  neutral  reaction,  and  are  either  fermentable  or  can,  in  most 
cases,  be  converted  into  substances  capable  of  fermentation.  Many 
of  the  carbohydrates,  especially  glucose,  are  easily  oxidized;  i.e.,  they  are 
good  reducing  agents,  as  is  shown  in  the  foregoing  experiments  by  the 
reduction  in  alkaline  solution  of  the  salts  of  copper  (Fehling's  solution). 

The  carbohydrates  are  now  conveniently  arranged  under  three  heads. 
They  are: 

1.  Monosaccharides  or  simple  sugars. 

This  group  contains  those  sugars  which  cannot  be  broken  down  into 
two  or  more  sugars.  Examples  of  these  are  glucose  and  fructose. 

2.  Polysaccharides  or  complex  sugars. 

Examples  are  cane  sugar,  sugar  of  milk  and  maltose. 

3.  Polysaccharides,  not  resembling  sugars. 
Examples  are  cellulose  and  starch. 

Glucose. — C6H12O6  (grape  sugar,  dextrose). — This  substance  is  widely 
distributed  throughout  the  vegetable  kingdom,  and  is  generally  accompa- 
nied by  fructose  or  fruit  sugar.  It  is  found  in  relatively  large  quantities 
in  the  juices  of  many  fruits,  such  as  the  grape,  fig,  strawberry,  cherry, 
mulberry,  etc.,  and  in  honey.  Traces  (o.i  per  cent,  or  less)  of  grape  sugar 
occur  in  the  liver,  in  the  normal  blood  and  in  the  urine.  In  the  disease 
diabetes  mellitus,  the  quantity  secreted  in  the  urine  is  greatly  increased, 
frequently  rising  as  high  as  5  to  10  per  cent.  Glucose  is  easily  prepared 
from  pure  cane  sugar,  which  is  easily  hydrolyzed  by  the  action  of  dilute 
acids,  yielding  glucose  and  fructose  as  indicated  by  following  equation: 

C12H220U  +  H20  ->  C6H1206  +  C6H1206 
Cane  sugar  Glucose        Fructose 

Commercially,  it  is  manufactured  in  large  quantities  by  heating  starch 
with  dilute  sulphuric  acid,  which  converts  the  starch  first  into  dextrin  and 
maltose,  and  then  into  glucose.  The  solution  obtained  by  this  process 
is  treated  with  chalk  to  remove  the  excess  of  acid,  and  filtered.  The 
filtered  solution  is  either  evaporated  to  a  syrupy  consistency  and  sold  as 
"glucose,"  or  to  dryness,  the  solid  product  obtained  being  known  as 
commercial  "grape  sugar." 


268  EXPERIMENTAL  CHEMISTRY. 

Glucose  is  soluble  in  water  and  alcohol.  Its  sweetness,  compared  with 
that  of  cane  sugar,  is  as  3  to  5.  It  is  fermentable  and  reduces  Fehling's 
solution.  It  is  optically  active,  turning  the  plane  of  polarization  to  the 
right.  (a)D  =  +  52.7°  for  a  10  per  cent,  aqueous  solution  which  has 
been  prepared  for  some  time.  Cold,  freshly  prepared  solutions  give 
a  much  larger  rotation. 

Saccharose,  C12H22On  (cane-sugar,  beet-sugar). — Sucrose  (saccharose) 
is  formed  in  the  juices  of  many  plants,  but  it  is  usually  associated  with 
substances  which  render  its  extraction  exceedingly  difficult  and  unprofit- 
able. The  sugar-cane,  sugar-beet,  sugar-maple  and  the  date-palm  are 
the  commercially  important  sources  of  sugar,  the  first  two  supplying  nearly 
all  the  sucrose  of  commerce. 

Sucrose  forms  white,  hard,  crystalline  granules,  but  may  be  obtained 
by  crystallization  from  water  in  large  monoclinic  prisms  (rock-candy). 
It  is  sparingly  soluble  in  alcohol,  but  dissolves  in  0.5  its  own  weight  of 
water  at  temperatures  below  20°  C.,  and  in  0.2  parts  at  100°  C.;  it  melts 
at  160-161°;  and  on  cooling  it  solidifies  to  a  pale-yellow,  amorphous, 
glassy  mass  known  as  barley  sugar,  which  after  a  long  time  becomes 
opaque  and  crystalline;  when  heated  to  200-210°  C.,  it  loses  water  and  is 
gradually  converted  into  a  brown,  almost  tasteless  substance  called 
caramel,  which  is  used  largely  for  coloring  liquors,  soups,  etc.  Oxidiz- 
ing agents  act  energetically  upon  cane  sugar,  converting  it  into  oxalic 
acid  and  saccharic  acid;  a  mixture  of  cane  sugar  and  potassium  chlorate 
will  deflagrate  with  explosive  violence  when  moistened  with  a  drop  of 
sulphuric  acid;  an  acid  solution  of  potassium  permanganate  is  decolorized 
(reduced)  by  sugar.  Although  sucrose  is  a  strong  reducing  agent,  it  does 
not  reduce  Fehling's  solution. 

Pure  sucrose  does  not  ferment,  but  if  a  trace  of  a  mineral  acid  is  added 
to  an  aqueous  solution  and  the  liquid  warmed  or  allowed  to  stand,  the 
sucrose  is  hydrolyzed  with  the  formation  of  equal  quantities  of  the  fer- 
mentable substances,  glucose  and  fructose. 

CBHBOil  +  H20-  C6HI206  +  C6H1206. 

Sucrose  Glucose        Fructose. 

Sucrose  is  optically  active,  (a)D  =  +  66.51°  at  20°  C.  for  a  solution 
containing  26.004  grams  of  sugar  in  a  volume  of  100  cm.3  Therefore, 
sucrose  is  dextrorotatory,  i.e.,  rotates  the  plane  of  polarized  light  to  the 
right,  but  when  it  has  been  hydrolyzed  with  acids  as  indicated  by  fore- 
going equation,  the  resulting  solution  of  glucose  and  fructose  is  levorota- 
tory,  i.e.,  the  plane  of  polarization  has  been  reversed  or  "inverted." 
This  is  due  to  the  fact  that  fructose*  causes  a  greater  rotation  to  the  left 
than  glucose  does  to  the  right.  The  mixture  of  glucose  and  fructose  is 
called  invert  sugar,  and  the  process  of  converting  sucrose  into  these  two 
forms  is  called  inversion.  An  instrument  known  as  the  saccharimeter 
is  used  to  determine  the  rotatory  power  of  a  known  solution  before  and 

*  (a)D=   — 93°  in  10  per  cent,  aqueous  solution. 


SOME  COMMON  CARBON  COMPOUNDS.  269 

after  inversion,  and  from  this  data  the  percentages  of  cane  sugar  and  in- 
vert sugar  present  in  the  sample  can  be  calculated. 

The  ease  with  which  sucrose  is  inverted  indicates  that  the  radicals  of 
glucose  and  fructose,  which  it  evidently  contains,  are  united  by  oxygen. 
(See  probable  formula  of  sucrose,  page  251.) 

Cellulose  (C6H10O5)n. — Cellulose  is  the  principal  constituent  of  cell 
membranes  and  of  the  woody  fiber  of  plants,  and  constitutes,  in  fact,  the 
framework  of  all  vegetable  tissues.  Some  parts  of  vegetables,  such  as 
cotton,  flax  and  hemp,  consist  almost  wholly  of  cellulose. 

Pure  cellulose  is  a  white,  transparent  mass,  insoluble  in  all  the  common 
solvents;  it  dissolves  slowly  in  concentrated  sulphuric  acid,  and  if  the  solu- 
tion be  diluted  with  water  and  boiled,  dextrine  and  finally  glucose  are 
produced.  Hence,  it  is  possible  to  convert  wood  into  sugar. 

Cellulose  is  used  extensively  in  the  manufacture  of  fabrics,  paper  and 
explosives. 

Paper  consists  essentially  of  cellulose  fibers  matted  or  felted  into  a  coher- 
ent layer  or  sheet.  Usually  it  is  "loaded"  with  mineral  matter  to  give 
it  weight  and  render  it  smooth  and  less  porous. 

Starch  (C6H10O5)n. — Starch  is  found  in  abundance  in  the  vegetable 
kingdom,  occurring  in  nearly  all  plants  in  a  greater  or  lesser  quantity. 
It  is  found  chiefly  in  the  seeds,  particularly  of  the  cereals  and  the  legumin- 
os&,  but  also  in  the  roots  and  stems  of  many  plants.  It  is  formed  in  the 
leaves  and  green  parts  and  then  transported  in  soluble  form  to  other  parts 
of  the  plant  where  it  is  used  to  build  up  the  tissues,  or  is  deposited  as  re- 
serve material.  The  details  of  the  formation  of  starch  are  not  fully 
understood,  but  it  appears  that  chlorophyl  acts  like  a  "contact"  substance 
or  a  "catalyser"  in  the  reaction  by  which  starch  is  formed  in  the  plant. 
The  carbon  dioxide  of  the  air  is  reduced  by  the  joint  action  of  sunlight 
(energy),  and  protoplasm  and  the  chlorophyl,  the  carbon  being  assimilated, 
while  a  portion  of  the  oxygen  is  set  free.  If  the  various  stages  in  which 
the  reaction  takes  place  are  disregarded,  the  ultimate  chemical  change 
may  be  represented  by  the  following  thermo-chemical  equation: 

6C02  +  5H20  —  C6H1005  +  602— 671,000  cal. 
or 

6CO2  +  6H2O  ->  C6H12O6  +  6O2. 

That  substances  of  such  a  high  molecular  weight  as  starch  or  glucose 
should  be  formed  directly,  as  indicated  by  above  equations,  is  regarded 
by  chemists  as  highly  improbable.  Baeyer*  holds  that  formaldehyde, 
CH2O,  is  the  first  reduction  product  of  carbon  dioxide  in  plants,  and 
that  glucose  from  which  starch  may  be  derived  is  formed  by  a  polymeri- 
zation process: 

6CO2  +  6H2O  —  6CH2O  +  6O2 
6CH20  -*  C6H1206 
C6H1206  -  C6H1005  +  H20. 

*  Berichte  der  deutchen  chcmischen  Gesellschaft,  3,  67. 


270  EXPERIMENTAL  CHEMISTRY. 

Starch  as  found  in  its  natural  condition  possesses  a  granular  structure, 
the  granules  differing  from  each  other  so  markedly  that  the  source  from 
which  they  are  derived  can  usually  be  determined  by  a  microscopic  exam- 
ination. Each  starch  granule  is  made  up  of  different  layers,  arranged 
abound  a  nucleus  which  is  generally  at  one  side  of  the  granule.  The  in- 
terior substance  of  the  granule  consists  of  "granulose,"  while  the  trans- 
parent, inert  and  insoluble  covering  resembles  cellulose*  in  structure. 

Ordinary  starch  is  a  white,  amorphous,  tasteless  substance,  insoluble 
in  water  and  alcohol,  but  if  heated  with  water  the  granules  swell  and  rup- 
ture the  cell  walls,  when  the  "granulose"  combines  with  water  to  form 
a  jelly-like  paste.  If  this  paste  is  boiled  with  an  excess  of  water  it  forms  a 
partial  solution  or  an  emulsion,  which  passes  readily  through  a  filter 
paper.  This  solution  yields  a  characteristic  blue  color  with  iodine,  hence 
the  use  of  the  latter  as  an  "indicator"  in  many  analytical  operations. 
Starch  is  levorotatory  and  is  converted  into  dextrin  or  British  gum  by 
heating  it  in  the  dry  condition  to  about  200°  C.  When  boiled  with  dilute 
acids,  starch  is  converted  successively  into  dextrin,  maltose  and  glucose. 
A  process  which  is  of  the  utmost  importance  in  the  manufacture  of  alcohol 
and  spiritous  liquors  is  the  decomposition  of  starch  at  about  65  °C.  by 
the  ferment  diastase  into  dextrin  and  maltose: 

3C6H1005  +  H20  —  C6H1005  +  C12H220U. 

Starch  is  an  important  article  of  food,  especially  when  associated  with 
albuminous  substances,  as  in  ordinary  flour. 

In  the  United  States  starch  is  manufactured  principally  from  maize 
and  wheat;  in  Europe,  mainly  from  potatoes,  but  also  from  maize,  wheat 
and  rice. 

*  "Recent  investigations  tend  to  prove  that  "starch  cellulose"  is  not  present  as 
such  in  the  granule,  but  is  formed  from  starch  substance  by  the  action  of  acids  or 
fermentation . " — Thorpe. 


CHAPTER  XXII. 
ALUMINUM   FAMILY. 

Boron,  B,  n  .  o 
(Aluminum,  Al,  27.1) 
(Earth  metals,) 

On  inspecting  the  periodic  table  it  will  be  seen  that  the  fourth  column 
contains  the  elements,  boron,  aluminum  and  a  number  of  rare  "  earth 
metals."  The  two  which  are  of  particular  interest  to  us  in  this  work  and 
are  of  the  most  importance  are  boron  and  aluminum.  The  relations 
of  these  two  are  very  similar  to  those  existing  between  carbon  and  silicon. 
While  boron  is  usually  classed  with  carbon  and  silicon  on  account  of  its 
close  resemblance  to  these  elements,  the  connection  between  boron  and 
aluminum  is  plainly  revealed  in  the  similarity  of  their  physical  and 
chemical  properties.  Boron  possessing  the  lower  atomic  weight,  the 
metallic  character  is  reduced  or  does  not  appear.  It  is  a  pronounced 
non-metal,  and  its  oxide  and  hydroxide  are  almost  exclusively  acidic. 
Its  properties  approximate  those  of  the  metalloids.  For  this  reason  it  is 
customary  to  treat  boron  with  them,  as  suggested  above. 

Aluminum  is  really  a  perfect  metal,  and  with  its  oxide  and  hydroxide 
the  basic  properties  predominate.  Although  the  hydroxide  is  capable 
of  forming  metallic  salts  with  strong  bases  (alkalies)  of  the  type  of  sodium 
aluminate  Na3AlO3,  yet  owing  to  the  higher  atomic  weight  of  aluminum 
the  basic  character  exceeds  the  acidic.  All  the  members  of  the  aluminum 
family  are  trivalent. 

Boron  will  be  treated  at  this  time  while  aluminum  will  be  taken  up  in 
connection  with  the  metals. 

BORON,  B. 
At.  Wt.  n.o       Sp.  gr.   2.5-2.63. 

As  may  be  inferred  from  that  which  has  preceded,  boron  occupies  a 
rather  isolated  position  among  the  metalloids,  and  may  best  be  regarded 
as  a  transition  element  between  this  class  and  the  metals.  Although  it 
seems  expedient  to  classify  boron  with  carbon  and  silicon,  the  elements 
which  it  resembles,  especially  in  the  free  state,  the  student  is  reminded 
that  the  group  of  elements  most  nearly  related  to  it  must  be  sought  for 
among  the  "earth  metals." 

The  element  boron  has  never  been  found  in  nature  in  the  free  state. 
It  occurs  most  abundantly  in  combination  as  boric  acid  and  as  metallic 

271 


272  EXPERIMENTAL  CHEMISTRY. 

borates,  among  which  are  tincal  or  crude  borax,  Na^E^C^,  boracite  (mag- 
nesium salt)  and  colemanite  or  borate  spar,  Ca^BgC^. 

Boron  is  an  infusible  solid  substance  capable  of  existing  in  both  the 
amorphous  and  crystalline  forms.  The  former  is  a  greenish-black  powder 
which  may  be  prepared  by  heating  the  oxide  of  boron,  B2O3,  with  magne- 
sium or  sodium  in  a  covered  crucible.  The  fused  mass  is  then  boiled 
with  hydrochloric  acid.  The  boron  is  separated  by  nitration.  Crystal- 
line boron  is  obtained  by  the  fusion  of  aluminum  with  boron  trioxide. 
On  account  of  its  hardness,  it  is  called  " adamantine  boron."  When 
prepared  by  the  above  process,  it  is  not  quite  pure.  It  contains  a  small 
quantity  of  aluminum  which  is  probably  isomorphous  in  this  form  with 
boron.  The  crystals  are  transparent,  and  in  their  luster  and  hardness 
they  resemble  the  diamond.  The  amorphous  variety  is  probably  unstable 
with  respect  to  the  crystalline.  The  former  has  a  sp.  gr.  of  2.5;  the 
latter,  2.63. 

Boron  forms  an  unstable  hydride  (probably  B3H3)  which  resembles 
stibine  in  its  readiness  to  undergo  decomposition.  However,  the  two 
most  important  compounds  of  boron  are  boron  trioxide  and  boric  acid. 
Although  the  acid,  H3BO3,  is  known  in  the  free  state,  the  salts  of  the  acid 
are  not  known  with  any  dergee  of  certainty.  The  salts  are  formed  from 
the  so-called  "  condensed  "  acids,  which  are  intermediate  products  formed 
by  the  dehydration  of  orthoboric  acid.  This  acid  looses  water  when 
heated  and  passes  into  the  anhydride,  B2O3,  which  melts  forming  a  glass- 
like  mass.  This  fused  substance  is  capable  of  dissolving  the  oxides  of 
various  metals,  many  of  which  impart  characteristic  colors  to  the  "  borax 
bead."  These  phenomena  serve  for  the  detection  of  such  metals  in  analy- 
sis. The  readily  fusible  alkali  salts  of  boric  acid  are  frequently  employed 
for  this  purpose,  and  for  cleaning  surfaces  to  be  soldered. 

Boric  acid  is  a  very  weak  acid,  the  salts  of  which  readily  undergo  hy- 
drolysis when  dissolved  in  water.  An  aqueous  solution  of  the  acid  con- 
ducts electricity  but  slightly  better  than  pure  water. 

Alcohol  and  boric  acid  interact  forming  the  corresponding  ester  which 
is  very  volatile.  If  the  alcohol  is  set  on  fire  the  ester  imparts  a  green  color 
to  the  flame.  This  reaction  is  frequently  used  for  the  detection  of  boric 
acid. 

Boric  acid  (boracic  acid)  is  used  widely  as  an  antiseptic  and  a  as  pre- 
servative. 

Experiment  I. — Preparation  and  Properties  of  Boron. 

Preparation  and  properties  similar  to  silicon.  Repeat  Exp.  I,  "  Silicon," 
substituting  boron  for  silicon,  and  boric  oxide  for  sand  or  quartz.  What 
is  the  valency  of  boron  ? 

Experiment  II. — Hydrolysis  of  Borates.  Boric  Acid.  Dehydration  of 
Boric  Acid.  Borax. 

(a)  Dissolve  a  little  powdered  borax  (sodium  tetraborate,  Na2B4O7) 


ALUMINUM    FAMILY.  273 

in  distilled  water.     Test  the  solution  with  neutral  litmus  paper.     Result? 
Explain.     Equation? 

(b)  Dissolve  20  grams  of  borax  in  100  cm. 3  of  hot  water.     Add  concen- 
trated hydrochloric  acid  until  the  solution  is  strongly  acid,  then  set  it 
aside  to  cool.     Results?     Equation? 

Filter  and  wash  out  the  mother-liquor  from  the  crystals  with  a  little 
cold  water.  Dry  some  of  the  boric  acid  crystals  by  pressing  them  be- 
tween filter  paper.  Test  the  solubility  of  the  crystals  in  cold  water  (?). 
In  alcohol  (?).  In  hot  water  (?).  Test  this  latter  solution  with  litmus 
paper  ( ?).  Determine  its  electrical  conductivity  relative  to  water.  What 
is  its  degree  of  ionization  ?  (See  tables.)  What  are  your  conclusions  as 
to  the  nature  of  boric  acid,  H3BO3?  Write  the  structural  formula  for 
H3B03. 

(c)  Place  about  a  gram  of  H3BO3in  a  small  porcelain  crucible;  heat 
gently  at  first,  then  strongly.     Results?     Equations  indicating  various 
stages  in  the  dehydration? 

(d)  Dissolve  5  grams  of  H3BO3  and  6  grams  of  Na2CO3  in  15  cm.3  of  hot 
water  (?).     Set  the  solution  aside  and  allow  it  to  evaporate  slowly? 
Result  ?    Equation  ?    Enumerate  some  of  the  uses  of  borax.    Its  formula  ? 

Experiment  III. — Preparation  and  Use  of  Borax  Beads.  Solubility  of 
Metallic  Oxides  in  the  Bead. 

(a)  Make  a  small  closed  loop  on  the  end  of  a  platinum  test  wire  which 
has  been  previously  sealed  into  a  piece  of  glass  tubing;  heat  the  wire  in 
the  flame  of  a  Bunsen  burner  and  dip  it  into  powdered  borax;  heat  it  in 
the  flame  until  the  white  puffy  mass  shrinks  to  a  small  glass-like  bead 
which  should  be  transparent.     If  the  bead  is  too  small  to  fill  the  loop, 
add  more  borax  and  heat  again.     The  bead  may  be  removed  by  dipping  it, 
while  hot,  into  cold  water;  the  sudden  cooling  shatters  it. 

Na2B4O7->  2NaBO2,  B2O3. 

(b)  Moisten  a  "borax  bead"  with  a  cobalt  chloride  or  nitrate  solution. 
Heat  the  bead  in  the  oxidizing  flame.     Notice  the  color  of  the  bead  when 
hot,  also  when  it  is  cold  (?).     The  color  is  usually  detected  by  looking 
at  the  bead  against  a  white  object  in  a  strong  light.     Is  the  color  of  the 
bead  altered  by  heating  it  in  the  reducing  flame? 

(c)  Repeat  (b),  substituting  a  copper  sulphate  solution  for  the  cobalt 
solution.     Tabulate  results.     Salts  of  the  following  may  also  be  used: 
nickel,  iron,  manganese  and  chromium. 

The  following  formulae  probably  represent  respectively  the  structures 
of  borax,  and  the  compound  which  is  formed  by  the  union  of  the  metallic 
oxide  and  the  fused  borax. 


18 


274  EXPERIMENTAL  CHEMISTRY. 

Na—  O—  B  Na—  O—  B 

/     \  /    \ 

00  00 

\  /  /      / 

B  Co  B 

\  \    . 

O  O 

B 

/     \ 
O         O 


O 

o 

/    \  I 

Co  B 


Na  —  O   —  B 


O     O 

\  / 
Na  —  O  —  B 


COLORS  OF  BORAX  BEADS. 


Metal. 

Chromium, 

Cobalt, 
Copper, 

Iron, 
Manganese, 

Nickel, 
Uranium, 


Oxidizing  Flame. 


Reducing  Flame. 


Cold. 

Oxide. 

Hot.               Cold. 

Green 

Cr203 

Emerald        Emerald 

green              green 

Blue 

CoO 

Blue                Blue 

Blue 

Cu2O 

Colorless         Red 

(opaque) 

Yellow 

Fe  O 

Bottle-green  Bottle-green 

Amethyst- 

MnO 

Colorless        Colorless 

red 

(Faint  rose)    (Faint  rose) 

Red-brown 

Ni 

Gray               Gray 

(Turbid)         (Turbid) 

U3Og         Green 


Green 


Experiment  IV. — Detection  of  Boron.     Tests. 


(a)  Dissolve  a  small  quantity  of  boric  acid  in  10  cm.3  of  alcohol  (ethyl 
or  methyl).     Pour  a  portion  of  the  mixture  into  an  inverted  porcelain 
crucible  cover  and  set  fire  to  it,  noting  the  color  of  the  flame  (?). 

(b)  To  a  strong  aqueous  solution  of  borax  in  a  test  tube  provided  with 
a  cork  carrying  a  glass  jet,  add  a  little  sulphuric  acid,  then  a  volume  of 
alcohol  equal  to  that  of  the  solution.     Heat  the  mixture  and  inflame  the 
vapor  issuing  from  the  jet  (?).     Equations? 

(c)  Dip  the  end  of  a  platinum  wire,  successively,  into  concentrated 
H2SO4,  glycerine,  and  powdered  borax.     Hold  the  wire  in  the  outer 
edge  near  the  bottom  of  a  small  Bunsen  flame.     What  color  is  imparted 
to  the  flame? 

(d)  Reaction  with  turmeric.     Dip  a  strip  of  turmeric  paper  into  a  solu- 
tion of  boric  acid,  or  a  borate  to  which  has  been  added  a  little  HC1  or 
H2SO4  to  liberate  the  boric  acid.     Observe  the  production  upon  the  paper 
of  a  characteristic  red-brown  stain.     The  color  is  distinguished  from  that 
produced  by  the  alkalies  by  the  fact  that  when  it  is  touched  with  a  drop  of 


ALUMINUM    FAMILY. 


275 


an  alkali  solution,  the  brown  color  is  changed  to  a  greenish-black  color, 
but  is  restored  to  its  original  color  and  is  not  discharged  by  dilute  HC1  or 
H2S04. 

The  only  borates  which  are  readily  soluble  in  water  are  those  of  the 
alkalies.  Magnesium  borate  together  with  a  few  others  are  difficultly 
soluble. 

The  similarity  of  boron  and  aluminum  may  be  inferred  from  the  data 
incorporated  in  the  appended  table: 


Physical  Properties. 
Atomic  weight, 
State  or  phase, 

Color, 

Specific  gravity, 

Specific  heat, 

Melting  point, 

Chemical  Properties, 

H-derivatives, 

State  or  phase, 

Halides, 

Heat  of  formation, 

State  or  phase, 

Stability, 

O-derivatives, 
Heat  of  formation, 

State  or  phase, 
Stability, 

Acids, 


Stability, 


Boron. 

II  .0 

Solid 

(Amorp. — Cryst.) 
Transparent  (Cryst.) 

2.5-2.63 


Aluminum. 
27.1 
Solid 

Silver-white 
2.58  (Hammered) 


Infusible 

BH3 

Gas 

BC13 

104,000  cal. 

Liquid 

Unstable  in  water 

B203 

317,200  cal. 

Solid 
Soluble 

(H3B03) 
(H  B02) 
(H2B407) 
Stable  in  water 


657° 


A1C13;A12C16 

32i,9oocal.  (A12C16) 

Solid 

Stable  in  water 

A12O3 
388,900  cal.(A!2O3, 

3H20) 
Solid 
Insoluble 

(H.A1  03) 
(H  Al  O2) 
(H2A1204) 
Unknown  in  the 
free  state. 


CHAPTER  XXIII. 
THE  METALS  OR  BASE-FORMING  ELEMENTS. 

Introductory  Note. — On  account  of  certain  physical  and  chemical 
properties  common  to  a  large  number  of  the  elements,  but  wanting  in  a 
greater  or  lesser  degree  in  others,  the  elements  are,  as  suggested  pre- 
viously, grouped  into  non-metals,  metalloids,  and  metals.  In  our  previous 
work  we  have  studied  chiefly  the  non-metallic  or  acid-forming  elements. 
We  are  now  ready  to  begin  the  systematic  study  of  the  metals,  having 
studied  but  one  typical  metal,  namely,  sodium. 

The  metals  have  already  been  defined  as  base-forming  elements.  They 
are  usually  good  conductors  of  heat  and  electricity,  and  are  endowed 
with  a  peculiar  luster  which  is  known  as  metallic  luster.  Although  the 
number  of  metallic  elements  is  much  larger  than  that  of  the  non-metals, 
the  chemistry  of  the  former  is  less  diverse  and  more  simple. 

The  student  is  again  reminded  that  there  is  no  sharp  distinction 
between  metals  and  non-metals,  but  rather  a  gradual  merging  of  one 
class  into  the  other,  depending  on  whether  the  distinction  is  based  upon 
their  physical  or  chemical  properties. 

Physical  Properties  of  the  Metals. 

1.  At   ordinary  temperatures  metals   are  solids  with  the  exception  of 
mercury. 

2.  The   relatively   high   conductivity  of    the  metals   for   electricity  is 
characteristic. 

3.  The  metals  when  in  the  compact  form  are  endowed  with  a  metallic 
luster.     With  the  exception  of  copper  and  gold,  the  metals  in  compact 
masses  possess  a  silver-white  color.     Most  of  the  metals  are  black  when 
in    a   powdered   condition,    magnesium    and    aluminum    being   notable 
exceptions. 

4.  Many  of   the    metals   possess  the  property  of  tenacity  in  a  very 
considerable  degree. 

5.  All  of  the  metals  can  be  obtained  in  the  crystallized  form. 

6.  The  metals  vary  widely  as  regards  specific  gravity  and  volatility 
and  fusibility.     (See   tables.)     Those   metals  which  have  a  sp.  gr.  less 
than  4  (5,  and  even  6,  are  sometimes  given  as  the  limit)  are  called  the 
light  metals,  and  those  with  a  greater  sp.  gr.,  the  heavy  metals. 

7.  Malleability  is  a  property  peculiar  to  most  of  the  metals.     Antimony 
and  bismuth  are  reduced  to  powder  when  hammered. 

8.  Many  of  the  metals  when  in  the  molten  condition  dissolve  in  one 
another  forming  mixtures  which  are  sometimes  spoken  of  as  undetermined 
compounds,  but  more  often  as  alloys.     These  alloys  frequently  possess 
properties  very  similar  to  solid  solutions.     They  are  generally  more 

276 


THE    METALS    OF    BASE-FORMING    ELEMENTS.  277 

fusible  than  their  component  metals.  "Wood's  metal"  is  an  alloy 
which  melts  at  about  66°  C.,  and  contains  bismuth  4  parts,  lead  2 
parts,  tin  and  cadmium,  each  i  part.  (See  Table  of  Alloys.)  Alloys 
in  which  mercury  is  one  of  the  components  are  known  as  amalgams. 

Chemical  Properties  Characteristic  of  the  Metals. 

1.  The   oxides  and   hydroxides  of   the  metals  usually  exhibit   basic 
properties. 

2.  The    metals    manifest    but   little    tendency  to  form  combinations 
with  hydrogen. 

3.  Each  metal  is  able  to  function  as  a  positive  radical  in  a  salt,  and  to 
exist  as  a  positive  ion  in  aqueous  solutions  of  said  salt. 

4.  The    metals    show  a   strong   attraction   for   the   halogen   group    of 
elements,  forming  compounds  which  as  a  rule  are  soluble  in  water  but 
are  riot  hydrolyzed. 

There  are  numerous  other  chemical  properties  which  are  characteristic 
of  metals,  and  which  may  serve  as  a  basis  of  distinction  between  metals 
and  non-metals,  but  they  are  not  common  to  all  the  metals.  The  student 
will  become  familiar  with  these  various  properties  as  the  work  progresses. 

Occurrence  oj  the  Metals  in  Nature. — Certain  metals  are  found  in  nature 
free  of  all  combination.  When  a  metal  cocurs  thus  it  is  said  to  be  native. 
Platinum,  gold,  silver,  copper,  bismuth,  arsenic,  antimony,  mercury,  etc., 
are  found  in  the  native  state.  More  often  the  metals  are  found  in  nature 
in  combination  with  one  or  more  of  the  other  elements,  as  oxides,  chlorides, 
sulphides,  carbonates,  sulphates,  phosphates  and  silicates.  When 
elements  or  compounds,  or  mixtures  of  elements  or  compounds,  possess- 
ing definite  physical  properties,  are  found  in  nature,  they  are  usually 
given  the  name  of  minerals.  Various  rocky  materials  known  as  the 
matrix  usually  accompany  the  minerals.  The  minerals,  or  the  minerals 
and  the  rocks  (matrix)  in  which  they  occur,  are  called  ores.  The  art  of 
extracting  the  metals  from  their  ores  is  called  metallurgy.  The  process 
of  extracting  the  metals  frequently  consists  of  heating  the  ore  with  a 
flux  which  combines  with  the  matrix  and  other  undesirable  portions, 
giving  a  fusible  slag.  The  separation  of  the  fused  slag  and  the  molten 
metal  is  easily  effected. 

The  following  outline  states  the  forms  in  which  the  metals  occur  in 
the  greatest  abundance,  and  are  of  chief  importance,  commercially: 

Native.  Carbonates. 

Gold  Barium 

Platinum  Calcium 

Silver  Copper 

Iron 

Chlorides  (chiefly).  Lead 

Potassium  Magnesium 

Sodium  Manganese 

Lithium  Strontium 


278  EXPERIMENTAL  CHEMISTRY. 

Silicates.  Arsenic 

Aluminum  Cadmium 

Calcium  Cobalt 

Lithium  Lead 

Magnesium  Mercury 

Potassium  Molybdenum 

Sodium  Nickel 

Silver 

Oxides.  Zinc 
Aluminum 

Copper  Sulphates. 

Chromium  Barium 

Iron  Calcium 

Manganese  Lead 

Tin  Strontium 
Zinc 

Phosphates. 

Sulphides  (chiefly).  Calcium 
Antimony 

Classification  oj  the  Metals. — There  has  been  much  discussion  as  to 
the  best  classification  of  the  metals.  The  question  cannot  be  answered 
categorically  and  it  is  well  to  remember  that  for  purposes  of  study,  the 
metals  may  be  grouped  variously,  according  to  the  selection  of  those 
properties  which  are  made  the  basis  of  comparison.  Valence  alone 
may  serve  for  classification,  and  in  such  case  the  arrangement  will  be 
very  similar  to  that  of  the  periodic  system.  Again,  the  scheme  of  classi- 
fication may  be  based  on  the  analytical  behavior  of  the  metals.  This 
latter  arrangement  will  bring  together  in  many  cases  those  metals  be- 
longing to  one  group  of  the  periodic  system,  but  in  a  few  instances  the 
elements  of  a  periodic  group  are  separated.  It  is  evident,  then,  that 
any  one  scheme  of  classification  must  be  one-sided  as  emphasizing  certain 
similarities  more  strongly  than  others. 

The  following  grouping  has  been  suggested  by  Ostwald: 

Non-metals.  Metals. 

Hydrogen  and  the  halogens.  Alkali  metals.  j  i^ght 

The  oxygen  group.  Alkaline  earth  metals.    > 

The  nitrogen  group.  Earth  metals. 

The  carbon  group.  The  iron  group.  j  fjeavy 

The  argon  group.  The  copper  group. 

Other  metals. 

In  the  work  which  follows  the  metals  will  be  considered  in  the  order 
in  which  they  occur  in  families  of  the  periodic  grouping.  This  arrange- 
ment affords  a  general  view  of  the  metals  as  brought  together  in  the 
natural  groups  of  elements: 


THE   METALS    OF    BASE-FORMING    ELEMENTS.  279 

Elements  oj  Group  I: 

Family  M.  (alkali  metals). — Lithium,  sodium  potassium,  rubid- 
ium, caesium  and  the  radical  ammonium,  NH4,  which  is  known  as 
the  "hypothetical"  metal. 

Family  m. — Copper,  silver  and  gold. 
Elements  oj  Group  II: 

Family  M  (alkaline  earth  metals). — Beryllium  (glucinum),  Mag- 
nesium, calcium,  strontium,  barium  and  radium. 

Family  m. — Zinc,  cadmium  and  mercury. 
Elements  oj  Group  III: 

Family  M  (earth  metals). — Scandium,  yttrium,  lanthanum, 
ytterbium. 

Family  m. — Boron,  aluminum,  gallium,  indium,  thallium. 
Elements  oj  Group  IV: 

Family  M. — Titanium,  zirconium,  cerium,  thorium. 

Family  m. — Carbon,  silicon,  germanium,  tin,  lead. 
Elements  oj  Group  V: 

Family  M. — Vanadium,  columbium,  and  tantalum. 

Family  m. — See  Nitrogen  Family. 
Elements  oj  Group  VI: 

Family  M. — Chromium,  molybdenum,  tungsten  and  uranium. 

Family  m. — See  Oxygen  Family. 
Elements  of  Group  VII: 

Family  M. — Manganese. 

Family  m. — See  Chlorine  Family. 
Transitional  Elements : 

Of  the  First  Long  Period. — Iron,  cobalt  and  nickel. 

Of  the  Second  Long  Period. — Ruthenium,  rhodium  and  pallad- 
ium. 

Of  the  Fourth  Long  Period. — Osmium,  iridium  and  platinum. 


CHAPTER  XXIV. 
ALKALI  METALS. 

Lithium,        Li,          7.03 
(Ammonium,  NH4,) 
Sodium,         Na,       23.05 
Potassium,     K,        39.15 
Rubidium,    Rb,      85.5 
Caesium,        Cs,      132.9 

The  metals  of  this  family,  together  with  their  corresponding  com- 
pounds, bear  a  very  close  resemblance  to  one  another.  This  family 
includes  those  metals  which  are  chemically  the  most  active.  The  activity 
increases  with  increased  atomic  weight.  The  metals  tarnish  very  quickly 
when  exposed  to  the  air  owing  to  rapid  oxidation.  They  decompose 


FIG.  46. 

water  violently,  liberating  hydrogen.  The  oxides  and  hydroxides  have 
strongly  basic  properties.  Their  salts  with  the  "active"  acids  do  not 
undergo  hydrolysis  when  in  aqueous  solution.  They  are  univalent,  and 
are  never  found  in  anions. 

The  compounds  of  ammonium  will  be  discussed  with  those  of  potas- 
sium to  which  they  show  a  marked  similarity. 

280 


ALKALI   METALS.  281 

LITHIUM,    Li. 

At.  Wt.  7.03         Sp.  Gr.  0.59. 
Experiment  I. — Flame  Color  of  Lithium  Compounds. 

Clean  a  platinum  wire  by  alternately  dipping  it  into  hydrochloric  acid 
and  heating  in  the  Bunsen  flame  until  it  gives  apparently  no  color  to  the 
flame,  then  dip  the  wire  into  a  dilute  solution  of  lithium  chloride  and  hold 
it  in  the  Bunsen  flame.  Notice  the  color  of  the  flame  and  examine  it 
with  the  spectroscope  (Figs.  46  and  47)  (Instructions)  (?).  Make  a 


FIG.  47. — Single  Prism  Spectroscope. 

diagram  of  the  "scale"  showing  the  relative  positions  of  the  sodium  or  D 
lines  (yellow)  and  the  lithium  lines.  Nearly  all  flames  in  the  laboratory 
show  the  yellow  lines  owing  to  the  presence  of  sodium. 

All  the  common  salts  are  readily  soluble  in  water  except  the  carbonate, 
phosphate  and  oxide  which  are  soluble  with  difficulty. 

SODIUM,  Na. 

At.  Wt.  23.05         Sp.  Gr.  0.97. 
Experiment  I. — Properties  of  Sodium. 

(a)  Recall  from  previous  work,  or  repeat  Exp.  I.,  "Sodium."     Read 
the  accompanying  note.     (See  "  Acids,  Bases  and  Salts.") 

1.  Tabulate  its  most  obvious  physical  properties. 

2.  Is   sodium   harder  or  softer  than  most  of  the  metals  with  which 
you  are  familiar? 


282  EXPERIMENTAL  CHEMISTRY. 

3.  Is  it  lighter  or  heavier  than  water? 

4.  Does  sodium  tarnish  when  exposed  to  the  air  ?     Explain. 

5.  Indicate    by  an    equation    the  interaction  of  water  and  metallic 
sodium. 

6.  Enumerate   the   chemical   and  physical  properties  which  indicate 
that  sodium  is  a  metal. 

7.  What  is  the  flame  color  of  sodium  compounds? 

(b)  Examine  at  least  ten  different  salts  of  sodium  (end  shelf),  noting 
the  obvious  physical  properties.  Give  the  name,  formula  and  color  of 
each  salt.  Determine  the  solubility  of  these  compounds.  Tabulate 
the  foregoing  data. 

Experiment  II. — Preparation  and  Properties  of  Sodium  Hydroxide. 

(a)  Dissolve  about  25  grams  of  soda  ash  (Na2CO3)  in  100  cm.3  to  125 
cm. ^  of  water  in  a  large  beaker,  and  heat  on  a  wire  gauze  to  boiling.     To 
10  grams  of  fresh  quicklime  which  has  been  powdered,  add  sufficient 
water  to  form  a  thin  paste — "milk  of  lime."     Use  heat  if  necessary  to 
start  the  slaking  action.    Add  the  milk  of  lime  gradually,  and  with  constant 
stirring  to  the  boiling  solution.     (Why?)     Continue  to  boil  for  several 
minutes  ( ?),  and  stir  constantly.     Remove  the  flame,  and  allow  the  pre- 
cipitate to  settle.     With  a  pipette  or  a  small  tube  remove  a  little  of  the 
liquid  to  a  test  tube,  and  test  for  the  presence  of  a  carbonate;  if  found 
to  be  present  add  more  milk  of  lime  and  boil;  but  if  absent,  decant 
the  liquid  into  a  bottle;  cork,  and  allow  it  to  stand  until  the  suspended 
solid  matter  settles.     Test  a  portion  of  the  precipitate  for  the  presence 
of  a  carbonate  ( ?).     Siphon  off  (or  filter  rapidly)  the  clear  supernatant 
liquid  by  means  of  a  glass  siphon  filled  with  water.     A  portion  of  the 
solution  may  be  evaporated  to  dryness  and  the  solid  substance;  the 
remaining  portion  should  be  preserved  in  a  tightly-stoppered  bottle  (?). 

(b)  Take  a  little  of  the  prepared  solution  between  the  fingers  (?). 
Test  it  with  litmus  paper  ( ?).     Try  the  flame  test  ( ?).     Add  an  excess 
of  the  solution  to  small  quantities  of  very  dilute  solutions  of  each  of  the 
following  substances  in  separate  test  tubes:  ferrous  sulphate  (?);  ferric 
chloride  ( ?) ;  zinc  sulphate  ( ?) ;  mercuric  chloride  ( ?) ;  copper  sulphate  (  ?). 
Boil  the  contents  of  the  tube  (?).     Describe  the  color  and  structure* 
of  each  of  the  precipitates  and  state  the  affect  of  boiling  same. 

(c)  Repeat   (b),    using   the   reagent  labeled   "NaOH  Solution"    (?). 
What  is  your  inference  as  to  the  identity  of  the  solution  prepared  in  (a)  ? 
By  what  other  names  is  sodium  hydroxide  known  ? 

(d)  Write  equations  representing  the  reactions  involved  in  (a)  and  (b). 

(e)  Examine  the  sodium  hydroxide  (solid)  as  found  on  the  side-shelf. 
(See  instructor.)     Expose  a  small  piece  to  the  air  on    a  watch   glass 
for  several  hours  (?).     Is  it  efflorescent  or  deliquescent? 

(/)  How  can  sodium  hydroxide  be  converted  into  sodium  chloride? 
Equation  ? 

*  Gelatinous,  flocculent,  curdy,  pulverulent,  granular  or  crystalline. 


ALKALI    METALS.  283 

Experiment  III. — Purification  of  Sodium  Chloride. 

Note. — Common  salt  usually  contains  such  impurities  as  calcium 
and  magnesium  chlorides,  sodium  sulphate,  calcium  sulphate,  etc. 
The  presence  of  the  chlorides,  magnesium  and  calcium,  causes  the  salt 
to  become  moist,  especially  in  damp  weather. 

Prepare  a  saturated  solution  of  sodium  chloride  by  grinding  50  grams 
of  salt  with  150  cm.3  of  water  in  a  mortar.  Filter  the  solution  into  a 
beaker  and  conduct  hydrogen  chloride  into  it.  It  is  suggested  that  this 
gas  be  prepared  by  placing  a  small  handful  of  salt  in  a  generating-flask, 
covering  it  with  concentrated  hydrochloric  acid,  and  allowing  concen- 
trated sulphuric  acid  to  drop  slowly  from  a  dropping-funnel  into  the 
mixture.  The  evolved  gas  is  passed  into  the  saturated  salt  solution  by 
means  of  an  inverted  funnel  or  thistle  tube  which  is  attached  to  the 
delivery  tube  of  the  generator.  The  mouth  of  the  funnel  should  dip 
just  below  the  surface  of  the  solution.  Pure  salt  separates  out  as  the 
operation  continues,  but  the  impurities  remain  in  solution.  (Why?) 
Explain  the  precipitation  of  the  salt  by  hydrogen  chloride.  Indicate 
the  action  by  means  of  "ionic"  equations.  When  the  precipitation 
has  continued  for  sometime  and  considerable  salt  has  separated,  remove 
the  generator,  and  allow  the  salt  to  settle,  then  decant  the  fluid  and  wash 
the  salt  with  15  or  20  cm.3  of  cold  water,  and  again  decant.  After  this 
process  of  washing  has  been  repeated  several  times,  the  last  traces  of 
water  may  be  removed  with  a  filter  pump  or  by  throwing  the  salt  upon 
filter  paper  and  pressing  the  salt  with  a  spatula.  Dry  the  salt  by  warming 
it  in  a  porcelain  dish  while  stirring  constantly  with  a  glass  rod. 

Describe  the  appearance  of  the  purified  salt.  Preserve  it  in  a  small 
bottle  for  future  use. 

Experiment  IV. — Preparation  of  Sodium  Carbonate  (Solvay  Process). 

To  80  cm.3  of  water  add  20  cm.3  of  ammonium  hydroxide.  Using 
the  foregoing  mixture  as  a  solvent,  prepare  a  saturated  solution  of  am- 
monium carbonate,  then  place  the  solution  in  a  corked  bottle  and 
saturate  it  with  sodium  chloride.  This  can  be  done  only  by  prolonged 
agitation  with  the  powdered  salt.  Allow  the  suspended  matter  to  settle, 
then  decant  the  clear  liquid  into  a  bottle  fitted  with  a  cork  and  a  delivery 
tube  which  reaches  to  the  bottom;  connect  the  delivery  tube  with  the 
laboratory  supply  of  carbon  dioxide  and  saturate  the  solution  with  the 
gas.  This  latter  operation  may  require  nearly  an  hour.  Finally  filter 
off  the  precipitated  matter,  and  dry  the  salt  by  pressing  it  between  filter 
papers.  Dissolve  a  little  of  the  salt  in  a  few  cm.3  of  water.  Dip  a  clean 
platinum  wire  into  the  solution  and  hold  it  in  the  colorless  Bunsen 
flame  (?).  Add  hydrochloric  acid  to  the  solution  and  test  the  evolved 
gas  for  carbon  dioxide  ( ?).  What  are  your  conclusions  as  to  the  identity 
of  the  salt  ? 

Experiment  V. — Purification  of  Sodium  Carbonate  by  Crystallization. 

Using  100  grams  of  sodium  carbonate,  prepare  a  saturated  solution  of 


284  EXPERIMENTAL  CHEMISTRY. 

the  salt  at  32°  C. ;  filter  at  the  same  temperature  and  collect  the  filtrate 
in  a  large  porcelain  dish;  cool  the  solution  to  o°  C.  by  packing  the  dish 
in  broken  ice;  heavy  layers  of  white  crystals  soon  form  in  the  porcelain 
dish.  Decant  the  supernatant  liquid  when  the  crystals  are  no  longer 
deposited,  then  dry  the  crystals  as  thoroughly  as  possible  by  pressing 
them  between  layers  of  filter  paper.  Redissolve  these  crystals  in  water 
at  33°  C.  and  repeat  the  foregoing  process.  The  crystals  are  subjected 
to  this  treatment  a  third  time,  after  which  they  are  pressed  and  dried. 
The  molecular  formula  of  the  salt  is  now  Na2CO3.  ioH2O.  Place  the 
crystals  between  watch  glasses,  and  allow  them  to  remain  exposed  to  the 
air  for  several  days.  They  effloresce,  and  in  a  comparatively  short 
time  are  transformed  into  white  powdery  Na2CO3.  Preserve  the  salt  in  a 
clean  glass-stoppered  bottle.  The  relative  solubility  of  sodium  carbonate 
at  various  temperatures  is  indicated  in  following  table: 

At  32°,  100  parts  of  water  dissolve  1140  parts. 
At  30°,  100  parts  of  water  dissolve  273  parts. 
At  20°,  100  parts  of  water  dissolve  92  parts. 
As  o°,  100  parts  of  water  dissolve  21  parts. 

Experiment  VI. — Dehydration  of  Hydrous  Sodium  Carbonate. 

Place  about  a  gram  of  hydrated  sodium  carbonate  in  a  test  tube  clamped 
so  that  its  mouth  is  inclined  slightly  downward;  heat  tube  to  a  dull  red- 
ness, and  note  the  change  in  the  salt.  What  is  evolved?  Dissolve  a 
small  quantity  of  the  anhydrous  salt  in  a  few  cm.3  of  water,  and  add  a 
few  drops  of  hydrochloric  acid  to  the  solution  ( ?).  Equations  ? 

Experiment  VII. — Action  of  Acids  upon  Sodium  Carbonate. 

Try  the  action  of  each  of  the  following  acids  upon  separate  portions 
of  a  solution  of  sodium  carbonate;  hydrochloric  acid  (?);  nitric  acid(?); 
sulphuric  acid  (?);  acetic  acid  (?).  Represent  the  reactions  by  ionic 
equations. 

Experiment  VIII. — Hydrolysis  of  Sodium  Carbonate. 

Test  a  solution  of  purified  sodium  carbonate  with  litmus  paper  (?). 
Explain  the  reaction  by  the  assistance  of  equations. 

Experiment  IX. —  Preparation  and  Decomposition  of  Sodium  Hydro- 
gen Carbonate. 

(a)  Preparation.     Make  a  saturated  solution  of  sodium  carbonate  by 
warming  (avoid  boiling)  it  with  water  in  a  beaker;  add  an  excess  of  the 
salt;  allow  the  solution  to  cool,  filter,  and  pass  carbon  dioxide  into  the 
filtrate  until  a  precipitate  forms  ( ?).     Indicate  the  reaction  by  an  equa- 
tion.    Is  sodium  hydrogen  carbonate  more  or  less  soluble  than  sodium 
carbonate  ?     Test  the  solution  with  litmus  paper  and  explain  its  reaction. 

(b)  Decomposition.     Place  about  i   gram  of  sodium  hydrogen  car- 
bonate (bicarbonate  of  soda)  in  a  test  tube,  add  10  cm.3  of  water  and 
heat  gently.     Test  the  evolved  gas  for  carbon  dioxide  (?).     Equations? 


ALKALI    METALS.  285 

Experiment  X.— Preparation  of  Sodium  Hydrogen  Sulphate. 
Recall  or  repeat  Exp.  X.  "Neutralization — Salts." 
Experiment  XI. — Reactions  of  Sodium  Salts. 

Note. — Use  the  purified  sodium  chloride  for  the  following  reactions. 

(a)  What  color  is  imparted  to  the  flame  by  sodium  salts?     Examine 
the  sodium  flame  by  means  of  a  spectroscope.     Locate  its  spectrum 
on  the  scale  (?). 

(b)  To  a  few  drops  of  a  strong  aqueous  solution  of  the  chloride  on 
a  watch  glass  add  five  or  six  drops  of  a  platinic  chloride,  PtCl4,  solution. 
Evaporate  carefully  to  small  volume.     Allow  the  mixture  to  cool,  and 
determine  the  solubility  of  the  red-colored  monoclinic  crystals  (?),  (a) 
in  water,  and  (b)  in  alcohol  (?). 

(c)  Sodium   cobaltic   nitrite,  Co(NO2)3.3NaNO2*,  produces   no  pre- 
cipitate in  acetic  acid  solutions  of  sodium  salts. 

(d)  Sodium  salts  are  but  slightly  volatilized  when  heated  in  a  porcelain 
dish  over  a  Bunsen  flame. 

All  sodium  salts  are  soluble. 

Experiment  XII.— (Quant.)     Alkalimetry  and  Acidimetry. 

(a)  Determination  of  the  strength  of  a  NaOH  solution  by  titration. 
Calculate  the  weight  of  pure  (99.95  per  cent,  purity)  oxalic  acid,  C2H2O4- 
.  2H2O,  required  to  make  500  cm.3  of  a  .  5N  oxalic  acid  solution.     Powder 
and  expose  on  a  watch-glass  for  20  min.  a  quantity  of  the  acid  slightly 
in  excess  of  the  calculated  amount,  then  weigh  out  on  glazed  paper  the 
exact  amount.     Transfer  the  oxalic  acid  to  a  500  cm.3  graduated  flask 
which  has  been  previously  calibrated,  and  fill  the  flask  to  the  mark  with 
distilled  water.     Fill  a  burette  with  the  solution.     Measure  carefully 
into  a  small  flask  from  another  burette  10  cm.3  of  the  sodium  hydroxide 
solution  prepared  in  Exp.  II,  dilutef  with  twice  its  volume  of  water,  and 
add  a  few  drops  of  a  phenolphthalein  solution.     Now  cautiously  run  acid 
into  the   flask  until  the  red  color  just  disappears.     Record  the  exact 
amount  of  acid  used,  and  calculate  (a)  the  weight  of  sodium  hydroxide 
per  liter,  (b)  the  normality  of  the  solution. 

(b)  If  the  sodium  hydroxide  solution  is  above  normal,  calculate  the 
volume    of    water  and    solution   necessary  for  the  preparation   of   100 
cm.3  of  a  normal  solution.     If  the  solution  is  less  than  normal  calculate 
the  amounts  required  for  250  cm.3  of  a  half-normal  solution  of  the  alkali. 

(c)  Using  the  amounts  calculated  in   (b)  , prepare  either  a  normal  or 
half-normal  solution  of  the  alkali. 

(d)  Using  the  solution  prepared  in  (c),  determine  the  normality  of  the 
*  Sodium  cobaltic  nitrite  may  be  prepared  as  follows:     To  33  grams  of  sodium  ni- 
trite add  100  cm.3  of  water,  make  slightly  acid  with  acetic  acid,  then  add  3.5  grams  of 
cobalt  nitrate.     Allow  the  solution  to  stand  for  several  hours,  and  filter  if  not  clear. 
The  solution  decomposes  slowly. 

f  A  concentrated  solution  may  decompose  the  indicator,  phenolphthalein. 


286  EXPERIMENTAL  CHEMISTRY. 

hydrochloric  acid  on  the  reagent  shelf.     Use  methyl-orange  (instructions) 
for  an  indicator. 

How  can  the  amount  of  acetic  acid  in  vinegar  be  determined  ? 

(e)  Define  alkalimetry.     Acidimetry? 

POTASSIUM,  K. 

At.  Wt.  39.15          Sp.  Gr.  0.87. 
Experiment    I. — Properties  of  Potassium. 

(a)  Observe  the  same  caution  in  the  manipulation  of  potassium  as 
suggested  in  the  use  of  sodium.     Place  a  piece  of  potassium  upon  a 
dry  paper  and  cut  off  a  piece  the  size  of   a  small  pea.     Observe  the  color 
and  luster  of  the  freshly  cut  surface  (?).     Note  the  effect  of  air  (?). 
Why  is  it  kept  under  kerosene  ?     Is  the  potassium  hard  ?     Half  fill  a 
tall  beaker  with  water   which  reacts  neutrally.     Scrape  the  coating  from 
the  piece  of   metal,  and  throw  it  (metal)  upon  the   water  contained  in 
beaker  which  should  be  covered  with  a  glass  plate.     Does  the  metal 
float  ?     Is  there  any  evidence  of  chemical  action  ?     Compare  its  interac- 
tion with  water  with  that  of  sodium  under  similar  conditions  ( ?). 

(b)  Wrap  a  piece  of  potassium  in  a  little  paper;  place  it  in  a  wire 
gauze  basket;  hold  it  under  water  and  collect  the  gas  which  escapes  in  a 
test  tube   by  displacement  of  water.     Apply  a  lighted   match   to   the 
mouth  of  the  tube.     Results?     Identify  the  gas  (?). 

(c)  Test  the  water  in  the  beaker  with  red  litmus  paper  (?).     Take  a 
little  of  the  water  between  the  fingers  (?).     Dip  a  clean  platinum  wire 
into  the  water  and  hold  it  in  the  colorless  Bunsen  flame.     Is  the  flame 
colored?     What  is  the  color? 

(d)  Pour  3   cm.3   of  the   shelf  reagent  labeled  potassium  hydroxide 
into  a  test  tube  and  dilute  with  an  equal  volume  of  water.     Repeat  (c), 
using     the     foregoing     solution.     Results?     Inference?     Indicate     the 
interaction  of  potassium  and  water  by  an  equation.     What  is  the  valency 
of  potassium? 

(e)  Place  a  drop  of  bromine  in  a  porcelain  dish  and  cautiously  drop 
a  fragment  of  potassium  into  it.     Results  ?     Equations  ? 

(/)  Put  a  small  piece  of  potassium  in  a  test  tube  with  a  little  iodine. 
Result  ?  Heat  the  tube  gently.  Results  ?  Equation  ? 

(g)  Enumerate  the  chemical  and  physical  properties  which  indicate 
that  potassium  is  a  metal. 

Experiment  II. — Potassium  Salts. 

Examine  at  least  ten  different  salts  of  potassium,  noting  the  obvious 
physical  properties.  Give  the  name,  formula,  color  and  solubility  of 
each  salt.  Tabulate  the  foregoing  data. 

Experiment  III. — Preparation  and  Properties  of  Potassium  Hydroxide. 

Proceed  as  in   Exp.  II,  "Sodium."     What  are  your  conclusions  as  to 


ALKALI    METALS.  287 

the   similarity   of    the  chemical   properties  of    sodium   hydroxide   and 
potassium  hydroxide? 

Experiment  IV. — Action  of  Acids  upon  Potassium  Hydroxide. 

(a)  Preparation  of  potassium  sulphate.    Dissolve  5  grams  of  potassium 
hydroxide  in  20  cm. 3  of  water.     Add  two  drops  of  a  phenolphthalein 
solution,  then  neutralize  the  solution  with  dilute  sulphuric  acid.     Evapo- 
rate until  crystallization  begins,    then    allow   the  solution  to  cool  and 
note  the  formation  of  crystals.     Equation  ? 

(b)  Preparation  of  potassium  hydrogen  sulphate.     Prepare  a  solution 
of  potassium  hydroxide  as  in  (a),  neutralize  with  a  measured  quantity  of 
dilute  sulphuric  acid,  then  add  a  volume  of  the  dilute  acid  equal  to  that 
used  in  neutralizing  the  alkali.     Evaporate  to  small  bulk  and  cool.      If 
a  crystalline  mass  is  not  obtained,  evaporate  further  and  cool.      Repeat 
if  necessary.     Equation  ? 

(c)  Repeat  (V),  using  nitric  instead  of  sulphuric  acid.     Equation? 

(d)  Repeat  (a) ,  substituting  hydrochloric  for  sulphuric  acid.     Equation  ? 

Experiment  V.— Hydrolysis  of  Potassium  Carbonate. 

Repeat  Exp.  VIII,  "Sodium,"  using  potassium  carbonate  instead  of 
sodium  carbonate. 

Experiment  VI.— Preparation  of  Potassium  Nitrate  by  Metathesis. 
Relative  Solubilities  of  Sodium  Chloride  and  Potassium  Nitrate. 

Dissolve  20  grams  of  sodium  nitrate  in  40  cm. 3  of  hot  water.  To  the 
boiling  solution  add  18  grams  of  potassium  chloride.  Stir  with  a  glass 
rod  until  all  the  salt  is  in  solution.  Evaporate  to  one-half  this  volume. 
What  separates  from  the  boiling  mixture?  Rapidly  filter  the  mixture 
by  decanting  the  hot  clear  liquid  from  the  deposit  of  salt  upon  a  filter. 
Set  the  filtrate  aside  to  be  examined  later.  Now  throw  the  deposited 
salt  upon  the  filter  and  press  out  the  mother  liquor  with  a  spatula. 
Recrystallize  this  salt.  Compare  these  crystals  with  any  that  may  have 
formed  in  the  filtrate  from  the  original  solution.  Identify  the  two  kinds 
of  crystals  (?). 

The  process  may  be  understood  more  clearly  by  inspecting  the  solu- 
bilities of  sodium  chloride  and  potassium  nitrate,  the  products  of  the 
interaction  of  the  factors.  The  following  table  gives  the  number  of 
grams  dissolved  by  100  cm. 3  of  water: 

Ato°C.         10°  C.  100°  C. 

Sodium  chloride,  35.6  grm.    35.63  grm.      39.9  grm. 

Potassium  nitrate,  13.3  grm.    25.00  grm.       247.0  grm. 

Experiment  VII. — Deflagration.     Gunpowder. 

(a)  Pulverize  separately  and  mix  intimately  3  grams  of  potassium 
nitrate  and  i  gram  of  sulphur.  Throw  the  mixture  into  a  red-hot  iron 
crucible  or  ignite  it  on  an  iron  plate.  Results?  Dissolve  the  residue 
in  water.  Devise  a  method  for  testing  for  potassium  sulphate  (?). 


288  EXPERIMENTAL  CHEMISTRY. 

(b)  Repeat  (a),  using  powdered  charcoal  instead  of  sulphur.     Test  the 
residue  for  potassium  carbonate  (?). 

(c)  Powder  separately  and  mix  thoroughly  on  paper  6  grams  of  potassium 
nitrate,  i  gram  of  sulphur  and  i  gram  of  charcoal.     Place  the  mixture  on 
an  iron  plate  or  brick  in  the  hood  and  ignite  it.     What  is  gunpowder? 
Define  deflagration. 

Experiment  VIII. — Preparation  of  Potassium  Cyanide  and  Potassium 
Thiocyanate. 

SeeExp.  XVI,  "Carbon." 

Experiment  IX. — Reactions  of  Potassium  Salts. 

Note. — Use  chemically  pure  potassium  chloride  for  the  following 
reactions: 

(a)  What  color  is  imparted  to  the  flame  by  potassium  salts  ?     Examine 
the  potassium  flame  by  means  of  a  spectroscope.     Locate  its  spectrum 
on  the  scale  with  reference  to  the  sodium  or  D  lines.     Make  a  diagram. 

(b)  To  a  few  drops  of  a  strong  aqueous  solution  of  the  chloride  on 
a  watch  glass  add  three  or  four  drops  of  a  solution  of  platinic  chloride, 
PtCl4.    Results?     Equation?     Determine  the  solubility*  of  the  yellow 
crystalline  precipitate  (a)  in  water  and  (b)  in  alcohol. 

(c)  Sodium  cobaltic  nitrite,  when  added  to  concentrated  solutions  of 
potassium  salts  which  contain  acetic  acid  but  no  free  inorganic  acid, 
gives  a  yellow  pulverulent  precipitate.     Equation? 

(d)  Potassium  salts  are  but  slightly  volatilized  when  heated  in  a  porce- 
lain dish  over  a  Bunsen  flame. 

Most  potassium  salts  are  soluble  in  water. 

AMMONIUM,    NH4. 

Experiment  I. — Preparation  and  Properties  of  Ammonium  Hydroxide. 

RecallExp.  II,  "Nitrogen."  Repeat  Exp.  II,  (b),  "Sodium."  Recall 
or  determine  the  relative  electrical  conductivities  of  normal  solutions  of 
potassium  hydroxide,  sodium  hydroxide  and  ammonium  hydroxide. 
(See  table,  Degree  of  lonization.)  What  ions  do  they  possess  in  common  ? 
What  are  your  conclusions  as  to  the  relative  activities  of  these  three  bases  ? 
Would  you  class  ammonium  hydroxide  among  the  active  bases?  Why? 
Can  ammonium  be  obtained  in  a  free  state?  In  what  state  or  condition 
does  it  exist  ?  What  name  is  frequently  applied  to  it  ?  Why  ? 

Experiment  II. — Preparation  of  Ammonium  Salts. 
Recall  Exp.  IV,  "Nitrogen." 

*  Potassium  chloro-platinate  (potassium  platinic  chloride)  is  soluble  in  no  parts 
of  water  at  10°  C.;  nearly  insoluble  in  alcohol,  and  quite  insoluble  in  a  mixture  of 
alcohol  and  ether. 


ALKALI    METALS.  289 

Experiment  III. —  Hydrolysis  of  Ammonium  Chloride. 

Test  a  solution  of  ammonium  chloride  with  litmus  paper.  Explain 
result.  Equation  ?  What  is  sal  ammoniac  ? 

Experiment  IV. — (Optional.)  Preparation  of  the  Unstable  Ammonium 
Amalgam. 

Place  about  50  cm.3  of  a  saturated  solution  of  ammonium  chloride  in 
a  dish  and  warm  gently.  Remove  the  flame  and  add  a  small  quantity 
of  sodium  amalgam  (see  Appendix)  to  the  solution  while  warm.  Results  ? 
Hold  a  piece  of  red  litmus  paper  over  the  dish  (?).  Equation  ? 

Hgx .  Na  +  NH4C1  —  Hgx .  NH4  +  NaCL 
2Hgx.NH4  -  2Hgx  +  2NHS  +  H2. 

Experiment  V. — Properties  of  Ammonium  Salts. 

(a)  Dissociation.     In  a  small  dry  evaporating  dish  heat  successively 
small  quantities  of  the  chloride,  the  nitrate  and  the  sulphate.     Note   the 
odor  of  the  fumes.     What  is  the  common  result  ?     Explain  the  formation 
of  the  white  cloud  which  appears  above  the  heated  salt.     Is  the  reaction 
reversible?     (Recall  Exp.  Ill,  "Chemical  Equilibrium.")     Equation? 

(b)  Decomposition.     Heat  a  small  quantity  of  ammonium  phosphate 
in  a  hard  glass  tube  (?).     Dissolve  the  residue  in  water  and  test  with 
litmus   paper.     Results?     Conclusions?     Equation?     Are    all   ammon- 
ium salts  completely  volatilized  by  simple  heating?     Is  ammonia  liber- 
ated from  all  salts  of  ammonium  when  they  are  heated  ?     (Recall  Exp.  I, 
"Nitrogen  and  the  Atmosphere,"  and  Exp.  XI,  (a),  "Nitrogen,"  before 
answering  this  last  question.) 

(c)  Purification  of  ammonium  chloride.     Plug  the  stem  of  a  large  glass 
funnel  with  clean  cotton-wool   and  invert  it  over  the  dish  in  which  a 
mixture  of  sand  and  ammonium  chloride  is  being  heated.     Continue  to 
heat  the  salt  until  a  heavy  white  layer  of  it  accumulates  oh  the  inside  of 
the  funnel.     Scrape  the  salt  on  to  a  piece  of  glazed  paper  and  examine 
it.     How  does  it  compare  in  purity  with  the  original  mixture  ?     Bottle 
the  salt  and  save  for  use  in   Exp.  VI. 

Exp.  VI. — Reactions  of  Ammonium  Salts. 

(a)  Do  ammonium  salts  impart  color  to  the  flame  ? 

(b)  Place  a  small  quantity  of  an  ammonium  salt  solution  in  a  beaker 
or  a  flask  and  add  a  few  cm.3  of   a  solution  of  sodium  hydroxide  or 
potassium  hydroxide.     Cover  the  beaker  with  a  watch  glass  on  the 
under  side  of  which  is  placed  a  moistened  piece  of  turmeric  paper  or 
red  litmus  paper.     Heat  the  beaker  gently.     Result.     Equation  ? 

(c)  To  a  few  drops  of  a  strong  aqueous  solution  of  the  chloride  on  a 
watch  glass  add  three  or  four  drops  of  a  solution  of  platinic  chloride. 
Result?     Equation?     Determine  the  solubility  of  the  yellow  crystalline 
precipitate  (a)  in  water  and  (b)  in  alcohol. 

This  yellow  precipitate  may  be  distinguished  from  the  similar  potas- 
19 


2QO  EXPERIMENTAL  CHEMISTRY. 

sium  salt  by  the  fact  that  when  heated  strongly  it  leaves  a  residue  of 
spongy  platinum. 

(d)  Sodium  cobaltic  nitrite  precipitates  from  acetic  solutions  of  am- 
monium salts  a  yellow  pulverulent  precipitate.     Equation  ? 

(e)  Nessler's  solution   (a  solution  of    potassium  mercuric  iodide  in 
potash)  is  used  to  detect  small  traces  of  ammonia  in  natural  waters. 
It  gives  a  brown  precipitate,  or  imparts  a  color  according  to  the  quantity 
of  ammonia  present. 

2(HgI2,2KI)  +  3KOH  +  NH3  —  NHg"2I  ,H2O  +  7KI  +  2H2O. 

(/)  Ammonium  salts  undergo  dissociation  or  are  decomposed  when 
heated  in  a  porcelain  dish. 

All  ammonium  salts  are  soluble  in  water.  Reagents  can  form  pre- 
cipitations only  in  concentrated  solutions. 

Experiment  VII. — Detection  of  the  Alkali  Metals  in  a  Mixture. 

How  can  the  salts  of  sodium  and  potassium  be  distinguished  from 
ammonium  salts  without  the  use  of  the  flame  and  the  spectroscope? 
How  could  you  distinguish  between  the  salts  of  sodium  and  potassium 
without  the  aid  of  the  flame  test  or  spectroscope?  Devise  a  system 
of  tests  which  will  provide  for  the  detection  of  the  three  alkali  metals 
in  a  solution  containing  their  respective  salts.  Apply  the  system  to  a 
" known"  solution  of  their  salts.  Ask  the  assistant  to  give  you  an 
"unknown"  solution.  Test  for  the  presence  of  the  alkali  metals.  Re- 
port (?).  Make  a  complete  record  of  all  \vork. 

RUBIDIUM,    Rb. 

At.  Wt.  85.5  Sp.  Gr.   1.52. 

See  text-book  and  lecture  notes. 

OESIUM,    CS. 

At.  Wt.   132.9          Sp.  Gr.   1.88 
See  text-book  and  lecture  notes. 

Lithium,  rubidium,  and  caesium  are  usually  placed  in  the  category 
of  rare  metals.  Lithium  salts  cost  about  10  s.  per  pound  while  rubidium 
and  caesium  salts  can  be  obtained  for  £  64  per  pound. 

PROBLEMS. 

1.  Calculate  the  percentage  of   sodium    sulphate  in    Glauber's  salt, 
Na2SO4.ioH2O. 

2.  What   weight   of    sodium  carbonate    is    theoretically  required   to 
prepare  100  grams  of  sodium  hydroxide? 

Na2CO3  +  Ca  (OH)a-»  2NaOH  +  CaCO3. 


ALKALI    METALS.  291 

3.  How    much    sal    soda   (washing   soda),   Na2CO3.  ioH2O,   can   be 
made  from  i  kg.  of  sodium  chloride?     The  equations  are: 

2NaCl  +  H2SO4  —  Na2SO4  +  2HC1. 

CaCO3  +  C  —  Na2CO3  +  CaS  +  4CO. 


4.  How  much  potassium  carbonate  is  necessary  to  prepare  100  grams 
of  potassium  hydroxide  ? 

5.  How  many  grams  of  sodium  hydroxide  are  necessary  to  liberate 
the  ammonia  from  50  grams  of  ammonium  nitrate  ?     From  50  grams 
of  ammonium  sulphate  ? 


CHAPTER  XXV. 

Copper,  Cu.  63  . 6 
Silver,  Ag.  107.93 
Gold,  AU.    197.2 

The  elements  of  this  family  occur  free  in  nature,  and  on  this  account 
are  among  those  metals  which  were  earliest  known.  They  are  used 
very  extensively  for  ornamental  purposes,  and  almost  universally  for 
coinage. 

These  metals  present  many  additional  properties  which  are  in  marked 
contrast  to  those  of  the  alkali  elements.  They  are  among  the  least 
active  metals.  They  do  not  interact  with  oxygen  or  water  at  ordinary 
temperature.  Silver  and  gold,  together  with  the  platinum  family,  are 
known  as  the  " noble  metals"  because  of  their  greater  resistance  to  the 
chemical  influences  of  air  and  water  and  their  easy  recovery  from  com- 
bination by  means  of  heat.  The  oxides  and  hydroxides,  with  the  ex- 
ception of  silver  oxide,  Ag2O,  possess  weakly  basic  properties.  The  hal- 
ides  of  these  metals,  save  those  of  silver,  are  hydrolyzed. 

Upon  the  basis  of  the  periodic  classification  these  three  elements  fall 
into  the  same  family,  but  in  many  of  their  chemical  relations  they  are  very 
dissimilar.  For  example,  silver  is  uniformly  univalent,  forming  one  series 
of  salts;  copper  is  both  univalent  and  bivalent,  forming  two  series;  gold 
is  univalent  and  trivalent,  and  likewise  forms  two  series.  Univalent 
copper  and  gold  bear  a  marked  resemblance  to  mercury,  but  bivalent 
copper  resembles  manganese  (Mn"),  zinc  (Zn"),  iron  (Fe")  and 
nickel  (Ni"),  and  trivalent  gold  bears  a  resemblance  to  aluminum  (Al"') 
and  iron  (Fe'").  On  the  other  hand,  silver  chloride,  AgCl,  and  cuprous 
chloride,  Cu2Cl2,  are  both  soluble  in  ammonia  but  insoluble  in  water. 

It  is  quite  obvious  that  that  family  is  not  homogeneous,  yet  in  many 
of  their  physical  attributes  these  metals  show  a  regular  gradation  in  their 
properties.  As  regards  ductility  and  malleability,  silver  is  intermediate 
between  copper  and  gold,  the  latter  possessing  these  properties  in  the 
maximum  degree.  With  respect  to  tenacity,  silver  is  also  intermediate, 
gold  being  the  least  tenacious  of  the  three  metals.  These  metals  are  the 
three  best  conductors  of  electricity.  The  electrical  conductivity  of  silver 
exceeds  that  of  all  other  metals. 

COPPER,    CU. 

At.  Wt.  63.6         Sp.  Gr.  8.95. 
Experiment  I. — Properties  of  Copper. 

(a)  Procure  a  piece  of  copper  wire,  and  clean  the  surface  by  scraping 
it  with  the  blade  of  an  old  pen-knife  or  by  scouring  it  with  sand.  Ob- 

292 


COPPER.  293 

serve  and  state  its  most  obvious  physical  properties.     Is  copper  a  good 
conductor  of  electricity?     Of  heat? 

(b)  Oxidation  of  copper.     Repeat  or  recall  Exp.  XVII,  (e),  "  Carbon." 
Equation  ?     What  change  is  slowly  undergone  by  copper  when  exposed 
to  the  air?     Upon  what  experimental  evidence  is  your  answer  based? 

(c)  Solution  tension.     Place  5  cm.3  of  a  solution  of  any  compound  of 
mercury  in  a  test  tube,  then  dip  a  clean  copper  wire  into  the  solution. 
After  a  short  time,  remove  the  wire  and  rub  it  carefully  with  a  soft  cloth. 
Result  ?     Has  any  of  the  copper  passed  into  solution  ?     Explain.     Equa- 
tion?    Compare  the  action  with  that  which  takes  place  when  zinc  acts 
upon  sulphuric  acid. 

Insert  an  iron  nail  or  wire  and  a  strip  of  zinc  into  separate  test  tubes 
half-filled  with  a  copper  sulphate  solution.  After  a  short  time  remove 
the  metals  and  examine  them.  Result?  Has  the  color  of  the  copper 
sulphate  solution  undergone  any  change  ?  Account  for  results  ?  Equa- 
tions ? 

(d)  Solubility  in  acids.     Using  small  portions  of  copper  filings  in  sepa- 
rate test  tubes,  determine  the  solubility  of  the  metal  in  hydrochloric  acid. 
In  nitric  acid  diluted  with  half  its  volume  of  water  ?     In  strong  sulphuric 
acid  ?     In  aqua  regia  ?     Use  heat,  if  necessary,  in  each  case  ?     Equations  ? 
What  color  is  common  to  the  various  solutions?     Recall  Exp.  IV,  (b},  (c). 
"The  Modern  Theory  of  Solution."     How  does  the  electrolytic  dissocia- 
tion theory  account  for  the  similarity  in  color  of  the  solutions  enumerated  ? 

(e)  Make  a  list  of  the  alloys  of  which  copper  is  one  of  the  components. 
(See  Appendix.) 

Experiment  II. — Reduction  of  Cupric  Oxide. 

(a)  Recall  Exp.  VIII,  "Hydrogen."     Equation? 

(b)  Thoroughly  mix  on  paper  about  2  grams  of  cupric  oxide  and  o.i 
gram  of  powdered  charcoal;   place  the  mixture  in  a  test  tube  and  heat. 
Test  the  evolved  gas  for  oxygen  and  carbon  dioxide,  and  examine  the 
residue  in  the  tube.     Results  ?     Equation  ? 

Experiment  III. — Preparation  and  Properties  of  Cuprous  Oxide. 

To  10  cm. 3  of  Fehling's  solution  in  a  test  tube  add  4  or  5  cm.3  of  a 
dilute  solution  of  glucose,  or  a  solution  of  cane  sugar  which  has  been 
boiled  with  a  few  drops  of  hydrochloric  acid  for  10  minutes.  (?).  Boil 
the  mixture*  vigorously  for  several  minutes.  The  red  precipitate  is  cu- 
prous oxide.  Explain  its  formation. 

Divide  the  precipitate  into  three  parts.  Expose  one  portion  to  the 
action  of  the  air  (?).  To  the  second  portion  add  strong  nitric  acid  and 
boil  (?).  Equation?  Add  ammonium  hydroxide  to  the  third  portion, 
and  shake  the  solution  in  a  test  tube.  Result? 

*  The  solution  must  be  strongly  alkaline. 


294  EXPERIMENTAL  CHEMISTRY. 

Experiment  IV. — Preparation  and  Properties  of  Cupric  Hydroxide. 
Cupric  Oxide. 

To  3  cm.3  of  a  solution  of  copper  sulphate  add  an  excess  of  a  solution 
of  sodium  hydroxide.  Results?  Write  the  reaction.  Boil  one-half 
the  mixture  and  account  for  the  change  in  the  color  of  the  precipitate. 
Equation  ?  To  the  other  half  of  the  mixture  add  ammonium  hydroxide. 
Result?  Equation?  Write  the  formula  of  the  complex  ion  which 
imparts  the  deep  blue  color  to  this  solution. 

Experiment  V. — Ammonio-cupric  Compounds,  Cuprammonium. 
Ammonio-cuprous  Compounds,  Cuprosammoniuni. 

(a)  Add  ammonium  hydroxide,  at  first  in  small  amounts,  then  in  ex- 
cess, to  a  solution  of  cupric  sulphate.     Results?     Give  the  formula  for 
ammonio-cupric  sulphate.     The  copper  now  forms  a  part  of  what  ion  ? 
Equations?     Would  you  infer  that  the  concentration  of  the  cupric  ion 
in  the  latter  solution  is  greater  or  less  than  in  the  copper  sulphate  solution  ? 

(b)  Place  10  or  15  cm.3  of  a  cuprammonium  sulphate  solution  in  a 
large  test  tube,  add  an  excess  of  copper  filings  or  turnings,  cork  air-tight, 
and  set  aside  until  the  solution  has  become  colorless.     Account  for  the 
change  in  color  of  the  solution  by  use  of  an  equation.     What  is  the  for- 
mula of  the  colorless  complex  ion?     Now  remove  the  cork  from  the 
tube  and  shake  the  solution  with  air.     Results  ?     Equation  ? 

Experiment  VI. — Preparation  and  Degree  of  Hydration  of  Cupric 
Chloride. 

(a)  Recall  Exp.  I,  (d).     Indicate  the  interaction  of  copper  and  hydro- 
chloric acid  by  an  equation. 

(b)  Put  2  to  3  grams  of  cupric  oxide  in  a  test  tube  or  small  flask,  add  a 
little  hydrochloric  acid  which  has  been  diluted  with  an  equal  volume 
of  water,  boil,  and  continue  to  add  acid  until  all  of  the  copper  oxide  is 
dissolved.     What  is  the  color  of  the  solution?     Equation? 

Place  three-fourths  of  the  solution  in  an  evaporating  dish,  evaporate, 
and  crystallize.  State  color  and  formula  of  crystals.  To  one-half  of 
the  remaining  solution  add  water  slowly  and  observe  the  color  (?).  To 
the  other  portion  add  concentrated  hydrochloric  acid  and  observe  the 
color.  Account  for  any  change  of  color.  Equations?  What  is  the 
color  of  anhydrous  cupric  chloride  ? 

(c)  Write  the  formulae  for  the  following  salts  of  copper:     Verdigris, 
blue  vitriol,  paris  green,  cupric  nitrate,  Scheele's  green,  copper  acetate. 

Experiment  VII. — Formation  and  Behavior  of  Cuprous  Chloride. 

(a)  To  3  grams  of  cupric  chloride  dissolved  in  15  cm.3  of  water  in  a 
small  flask,  add  5  cm.3  of  concentrated  hydrochloric  acid  and  about 
7  grams  of  copper  filings  or  turnings;  an  excess  of  copper  must  be  used. 
Boil  the  mixture  until  it  is  colorless,  or  a  drop  of  the  brownish-black 
liquid  gives  a  white  precipitate  when  added  to  a  test  tube  full  of  water. 


COPPER.  295 

Pour  the  liquid  into  a  large  volume  of  water  in  a  beaker.  Wash  the 
white  precipitate  by  decantation,  and  keep  it  under  water  until  used. 
What  change  have  the  cupric  ions  undergone  ?  Equation  ? 

(b)  Expose  a  little  of  the  cuprous  chloride  covered  with  water,  in 
a  test  tube,  to  the  sunlight.  Result  ?  The  reaction  is  said  to  be  repre- 
sented by  the  following  equation: 

l—  CuCl2  +  Cu. 


(c)  Cuprous    chloride    dissolves    in    concentrated    hydrochloric    acid 
giving  the  two  colorless  complex  acids,  HCuCl2  and  H2CuCl3.     Equa- 
tions?    What  is  the  complex  anion  formed  as  the  result  of  the  above 
reaction?     Do  complex  ions  tend  to  dissociate?     Show  by  ionic  equa- 
tions the  various  equilibria  involved  in  a  solution*  of  cuprous  chloride 
in  hydrochloric  acid. 

(d)  To  a  hydrochloric  solution  of  cuprous  chloride  add  a  little  concen- 
trated nitric  acid.     The  resulting  color  of  the  solution  shows  what  ion 
to  be  present?     Explain. 

Define    oxidation    in    terms   of    the    electrolytic    dissociation    theory. 

(e)  Cuprous  chloride  is  soluble  in  ammonium  hydroxide  giving  a  color- 
less solution  which  quickly  turns  to  a  deep-blue  owing  to  the  oxidizing 
action  of  the  air.     Equations? 

(/)  Precipitation  of  hydrated  cuprous  oxide.  Add  an  excess  of  a  sodium 
hydroxide  solution  to  a  hydrochloric  acid  solution  of  cuprous  chloride. 
Result?  Equation?  Divide  the  mixture  into  two  portions.  Boil  one 
portion.  Result?  Equation?  Shake  the  other  portion  with  air  in  a 
test  tube.  Result  ?  Equation  ? 

Experiment  VIII.  —  Precipitation  of  Cuprous  Iodide. 

Add  a  small  quantity  of  a  solution  of  potassium  iodide  to  10  cm.3 
of  a  very  dilute  solution  of  cupric  sulphate.  Observe  the  color  of  the 
solution  and  the  precipitate  (  ?).  Filter  and  wash  the  precipitate.  Divide 
the  nitrate  into  two  parts.  Add  one  part  to  a  dilute  cold  starch  emul- 
sion (?).  Shake  the  other  part  with  a  little  chloroform  or  ether  (?). 
Equations  ? 

Experiment  IX.  —  Dehydration  of  Hydrated  Cupric  Sulphate. 

See  Exp.  VIII,  (a),  "  Water."   Powder  a  small  crystal  of  cupric  sulphate, 
and  dehydrate  it  by  heating  it  gently  in  a  test  tube  clamped  in  a  nearly 
.  horizontal  position.     Results?     Equation?     What  is  the  effect  of  a  little 
water  upon  the  anhydrous  salt? 

Experiment  X.  —  Decomposition  of  Cupric  Nitrate.  Preparation  of 
Pure  Cupric  Oxide. 

Dissolve  5  grams  of  cupric  nitrate  in  25  cm.3  of  distilled  water,  filter  off 
any  insoluble  particles,  evaporate  the  filtrate  to  dryness  in  the  hood, 

*  The  solution  of  cuprous  chloride  in  hydrochloric  acid  or  ammonium  hydroxide  is 
a  valuable  reagent  used  in  gas  analysis  for  the  absorption  of  carbon  monoxide. 


296  EXPERIMENTAL  CHEMISTRY. 

and  heat  the  dry  residue  in  a  porcelain  crucible  until  it  is  of  a  uniform 
black  color.     Equation? 

Experiment  XI.  —  Precipitation  of  Basic  Cupric  Carbonate. 

To  a  solution  of  cupric  sulphate  add  a  slight  excess  of  a  sodium  car- 
bonate solution  (  ?).  Explain  the  participation  of  water  in  this  reaction. 
Boil  the  mixture  with  an  excess  of  the  precipitant  (?).  Equation?  What 
is  malachite?  Its  formula? 


2CuS04  +  2Na2C03  +  H20  <±  Cu2(OH)2CO3 

2Cu2(OH)2CO3  +±  Cu4O3(OH)2  +  CO2  +  H2O. 

Experiment  XII.  —  Precipitation  of  Cupric  Ferrocyanide. 

Add  a  solution  of  a  potassium  ferrocyanide  to  a  very  dilute  solution  of 
cupric  sulphate.  Result  ?  Equation  ? 

Experiment  XIII.  —  Preparation  of  a  Double  Salt.  Potassium-Cupric 
Sulphate. 

Using  5  grams  of  powdered  cupric  sulphate,  prepare  a  saturated  solution 
at  70°.  Calculate  the  weight  of  powdered  potassium  sulphate  which  must 
be  taken  to  give  the  same  fraction  of  its  molecular  weight.  Dissolve  this 
weight  of  salt  in  its  own  weight  of  water  at  70°,  and  add  a  few  drops  of 
sulphuric  acid.  Mix  the  two  solutions,  taking  care  to  secure  complete 
solution  of  both  salts  before  mixing.  Pour  the  mixture  into  a  beaker  or 
crystallizing  dish  and  set  it  aside  to  crystallize.  Examine  the  whitish- 
blue  crystals  and  compare  them  with  those  of  cupric  sulphate  (  ?).  Give 
the  molecular  formula  of  the  crystals.  Dry  the  crystals  between  layers 
of  filter  paper  and  preserve  them  in  a  stoppered  bottle  for  future  use. 

Dissolve  a  part  of  the  crystals  in  water.  Is  the  cupric  ion  present  in 
the  solution?  By  what  experimental  evidence  can  you  support  your 
answer  ? 

Do  "compound  salts"  yield  complex  ions? 

Experiment  XIV.  —  (Quant).     Equivalent  of  Copper. 

Dissolve  an  accurately  weighed  amount  (about  2  grams)  of  cupric  sul- 
phate in  distilled  water  in  a  beaker.  Record  the  weight.  Procure 
a  small  rod  or  bar  of  pure  zinc,  file  and  scour  it  until  its  surface  is 
smooth  and  clean;  wash  and  dry  it,  then  weigh  accurately.  Record  the 
weight.  Place  the  zinc  in  the  solution  and  allow  it  to  remain  there  until 
the  latter  becomes  colorless.  Take  the  zinc  out  of  the  solution,  carefully 
remove  the  brown  deposit  (  ?)  by  directing  a  small  stream  of  water  against 
it;  dry  and  weigh  it.  Record  the  weight.  What  weight  of  zinc  has  gone 
into  solution  ?  Assuming  that  all  of  the  copper  has  been  displaced  from 
the  solution,  calculate  from  the  formula  the  weight  of  copper  deposited 
upon  the  zinc.  This  weight  of  copper  will  be  the  equivalent  of  the 
weight  of  zinc  in  solution. 

The  atomic  weight  of  zinc  is  65.4.  Calculate  the  atomic  weight  of 
copper  from  above  data. 


COPPER.  297 

Experiment  XV. — (Quant.)  Determination  of  the  Percentage  of 
Copper  in  Cupric  Sulphate  Crystals  by  Electrolysis. 

Dissolve  an  accurately  weighed  amount  of  pure  cupric  sulphate 
in  water  in  a  weighed  platinum  crucible  or  dish.  Use  10  cm.3  of  water 
for  each  .1  gram  of  salt.  Add  10  drops  of  concentrated  nitric  acid  to 
make  the  solution  thoroughly  acid,  and  to  prevent  the  copper  from 
depositing  in  a  loose,  spongy  form  instead  of  a  hard,  smooth  plate.  Cover 
the  crucible  with  a  perforated  watch  glass,  and  suspend  a  flat  spiral 
of  platinum  in  the  solution.  Place  the  crucible  on  an  electrolysis  stand. 
(Instructions.)  Connect  the  stand  with  the  source  of  electricity  (a  bat- 
tery of  four  to  eight  gravity  cells  may  be  used)  in  such  a  manner  that  the 
platinum  spiral  becomes  the  anode  and  the  crucible  the  cathode.  The 
density  (ND100)  of  the  current  should  not  be  greater  than  i  ampere,  nor 
should  the  electrode  tension  exceed  2.5  volts.  As  the  current  continues 
to  act,  observe  the  changes  which  occur  in  the  crucible  (?).  To  deter- 
mine when  the  deposition  of  the  copper  is  complete,  a  drop  of  the  liquid 
is  taken  out  from  time  to  time  after  about  five  hours,  and  tested  on  a 
porcelain  surface  with  a  drop  of  potassium  ferrocyanide  which  has  been 
acidified  with  hydrochloric  acid.  When  the  electrolysis  is  finished,  the 
crucible  is  removed,  carefully  and  repeatedly  washed  with  water  and  then 
with  small  quantities  of  alcohol,  and  dried  in  an  air  bath  at  90°  C.  De- 
siccate and  weigh.  What  is  the  weight  of  the  copper  deposited  ?  Calcu- 
late the  percentage  of  copper  by  weight  in  cupric  sulphate  crystals.  Com- 
pare with  that  calculated  from  the  molecular  formula. 

Explain  electro-plating  in  terms  of  the  electrolytic  dissociation  theory. 

Experiment  XVI. — Analytical  Reactions  of  Cupric  Salts. 

(a)  Dip  a  platinum  loop  into  a  solution  of   any  cupric  salt  and  hold  it 
in  the  Bunsen  flame.     The  compound  imparts  to  the  flame  a  brilliant 
blue  color  which  instantly  changes  to  green  when  the  supply  of  acid  is 
exhausted. 

(b)  Mix  a  small  quantity  of  any  cupric  salt  with  sodium  carbonate  and 
heat  strongly  upon  charcoal  in  the  reducing  flame  (?).     Equation? 

(c)  The  borax  bead  test.     Copper  compounds  give  what  color  to  it 
(a)  when  heated  in  the  oxidizing  flame;     (b)  in  the  reducing  flame? 
Observe  in  each  case  whether  the  color  of  the  bead  while  hot  is  identical 
with  that  when  the  bead  is  cold. 

(d)  Prepare  a  dilute  solution  of  pure  cupric  sulphate  and  test  the  reac- 
tion of  the  solution  with  litmus  paper  (?).     Explain  the  result. 

Do  the  salts  of  active  metals  undergo  hydrolysis?  What  is  your  in- 
ference as  to  the  chemical  activity  of  copper  as  a  metallic  element  ? 

(e)  Pass  hydrogen  sulphide  through  5  cm.3  of  a  dilute  solution  of  cupric 
sulphate  (?).     Filter  and  divide  the  precipitate  into  three  parts. 

Determine  the  solubility  of  cupric  sulphide  in  sulphuric  acid  (?). 

Write  the  reaction  for  the  precipitation  of  cupric  sulphide  from  cupric 
sulphate  with  hydrogen  sulphide.  Is  the  reaction  reversible?  Upon 
what  experimental  evidence  do  you  base  your  answer? 


298  EXPERIMENTAL  CHEMISTRY. 

Ascertain  if  cupric  sulphide  is  soluble  in  (a)  nitric  acid  ( ?),  (b)  yellow 
ammonium  sulphide  (?).  Equations? 

(/)  To  a  dilute  solution  of  cupric  sulphate  add  a  potassium  cyanide 
solution  (poison)  drop  by  drop.  Observe  the  precipitation  of  the  un- 
stable cupric  cyanide,  which  readily  decomposes  into  cyanogen  and 
cuprous  cyanide.  Equations? 

Is  cuprous  cyanide  soluble  in  an  excess  of  potassium  cyanide  ?  Does 
the  resulting  solution  show  the  color  of  the  cupric  ion  ?  Is  the  cupric  ion 
present  ?  If  not,  of  what  ion  does  the  copper  now  form  a  part  ?  What 
general  name  is  given  to  this  class  of  ions?  Equation? 

Pass  hydrogen  sulphide  through  a  solution  of  cupric  sulphate  which 
has  been  decolorized  by  the  addition  of  an  excess  of  potassium  cyan- 
ide (?).  Do  results  indicate  the  absence  or  presence  of  the  cupric  ion? 
Explain. 

(g)  The  deep-blue  color  imparted  to  solutions  containing  copper  salts 
by  the  addition  of  an  excess  of  ammonium  hydroxide,  and  the  deposition 
of  copper  upon  strips  of  iron  or  zinc  when  immersed  in  solutions  of  salts 
of  the  former  are  tests  frequently  used.  Equations? 

Experiment  XVII. — Experimental  Study  of  Compound  Salts  and  Com- 
plex Salts. 

Using  the  crystals  preserved  from  Exp.  XIII,  prepare  a  solution  of 
potassium-cupric  sulphate.  Is  this  a  double  salt  or  a  complex  salt? 
Is  the  cupric  ion  present  in  the  solution  ?  Give  a  reason  for  your  answer. 
Pass  hydrogen  sulphide  through  the  solution  ( ?).  Do  the  results  support 
your  previous  conclusion  as  to  the  presence  of  the  cupric  ion  ?  Explain. 

Recall  or  repeat  Exp.  XVI,  (/). 

Explain  in  terms  of  the  ionic  hypothesis  the  difference  between  "  double 
salts"  and  "complex  salts." 

The  scheme  of  ionization  for  "double  salts"  is  identical  with  that  of 
"complex  salts." 

Double  salt,  K2SO4.CuSO4.6H2O: 

K2Cu(S04)2  <±  2K-  +  Cu(S04)"2  (I) 

Cu(SO4)"2<=>  CuSO4  +  SO"4  (II) 

Cu(SO4)    <±  Cu"  +  SO"4  (III) 

Complex  salt,   K4Fe(CN)6      +±  4K/  +  Fe(CN)""  (F) 

Fe(CN)""6  <z>  Fe(CN)2  +  4CN'  (IF) 

Fe(CN)2      <=»Fe"  +  2CN'  (IIF) 

The  ionization  corresponding  to  equations  (II)  and  (III)  are  practi- 
cally complete  for  double  salts  at  infinite  dilution,  and  for  complex  salts 
under  similar  conditions  the  reactions  corresponding  to  equations  (IF) 
and  (IIF)  take  place  only  to  a  very  limited  extent;  in  fact,  so  small,  that 
our  ordinary  chemical  tests  fail  to  indicate  the  presence  of  the  Fe"  ion. 


SILVER.  299 

The  results  of  refined  measurements  indicate  that  simple  ions  exist 
to  a  very  small  extent  in  dilute  solutions  of  complex  salts  and  that  com- 
plex ions  are  found  in  solutions  of  double  salts,  except  at  great  dilution. 

It  is  obvious  that  the  distinction  between  "double"  and  "complex" 
salts  is  one  of  degree  and  not  of  kind. 

SILVER,    Ag. 

At.  Wt.  107.93     Sp.  Gr.  10.5. 
Experiment  I. — Properties  of  Silver.     Preparation  of  Silver  Nitrate. 

(a)  Examine  a  small  piece  of  silver  and  make  a  record  of  the  most 
obvious  physical  properties  of  the  metal. 

Does  silver  undergo  oxidation  when  exposed  to  the  air?  Upon  what 
observation  do  you  base  your  answer? 

(b)  Dissolve  a  small  piece  of  silver  in  nitric  acid  diluted  with  an  equal 
volume  of  water  ( ?).     Evaporate  the  solution  to  small  bulk  and  crystallize. 
Dry  the  crystals  and  place  them  in  a  bottle.     Use  these  crystals  for  pre- 
paring solutions  of  silver  nitrate.     Indicate  the  interaction  of  nitric  acid 
and  silver  by  means  of  an  equation. 

Experiment  II. — Electrolytic  Deposition  of  Silver.     "Silver  Tree." 

Dissolve  about  15  grams  of  silver  nitrate  in  40  cm.3  of  water.  Place  the 
solution  in  a  small  glass  beaker  or  crystallizing  dish.  Introduce  platinum 
electrodes  into  the  solution  so  that  they  will  be  diametrically  opposite  to 
one  another.  Join  these  electrodes  to  the  battery  wires.  Allow  the  ac- 
tion to  continue  for  30  minutes  or  longer.  A  beautiful  deposit  of  silver 
will  be  made  on  the  cathode.  This  deposit  which  spreads  from  the 
cathode  toward  the  anode  is  called  the  "silver  tree"  because  of  its  strong 
resemblance  to  vegetable  growth. 

Experiment  III. — Replacement  of  Silver  in  its  Compounds  by  Copper. 

Place  a  strip  of  iron,  zinc,  tin,  lead  or  copper  in  a  solution  of  silver 
nitrate  contained  in  a  test  tube.  Result  ?  Ionic  equation  ?  Examine 
the  table  of  "Solution  Tensions."  Name  the  above  metals  in  the  order 
of  increasing  solution  tension.  Does  silver  ordinarily  displace  hydrogen 
from  aqueous  solutions  of  the  acids? 

Experiment  IV. — Precipitation  of  Silver  Oxide. 

To  a  dilute  solution  of  silver  nitrate  add  a  sodium  hydroxide  solu- 
tion ?  Result  ?  Equations  ?  Is  the  precipitate  soluble  in  excess  of  the 
precipitant  ? 

Experiment  V. — Preparation  and  Behavoir  of  the  Halides  of  Silver. 

(a)  Silver  chloride.  To  5  cm. 3  of  a  dilute  solution  of  silver  nitrate  add 
a  slight  excess  of  dilute  hydrochloric  acid  (?).  Equation?  Filter  and 
wash  the  precipitate  with  water. 


3OO  EXPERIMENTAL  CHEMISTRY. 

Put  a  small  part  of  the  precipitate  upon  a  watch  glass  and  expose  it  to 
direct  sunlight  (?).  Equation? 

To  a  second  portion  of  the  precipitate  add  ammonium  hydroxide  ( ?). 
Equation  ?  What  complex  ion  is  formed  ?  Now  add  dilute  nitric  acid 
drop  by  drop  in  excess  (?).  Equation? 

Treat  the  third  portion  of  the  precipitate  as  in  the  preceding  case,  using 
an  ammonium  carbonate  solution  instead  of  Ammonium  hydroxide. 
Equations  ? 

(b)  Silver  bromide.     Repeat  (a),  substituting  a  potassium  bromide  solu- 
tion for  the  hydrochloric  acid. 

(c)  Silver  iodide.     Repeat  (a),  using  a  solution  of  potassium  iodide  in 
place  of  the  hydrochloric  acid. 

What  extensive  practical  use  is  made  of  the  change  produced  in  silver 
salts  by  light  ?  Do  the  halides  of  silver  show  any  gradation  as  to  color  ? 
If  so  state  it? 

Experiment  VI. — Preparation  of  Pure  Silver  from  an  Alloy. 

Dissolve  a  copper-silver  alloy  in  dilute  nitric  acid  in  a  beaker  or  cas- 
serole. Use  heat  if  necessary.  Does  the  solution  possess  any  property 
which  indicates  the  presence  of  copper?  Explain.  Evaporate  just  to 
dryness,  and  take  up  the  residue  with  warm  water.  Filter,  add  a  few 
drops  of  nitric  acid  to  the  filtrate,  and  while  stirring  add  dilute  hydro- 
chloric acid  in  slight  excess.  Continue  to  stir  the  mixture  until  the  major 
portion  of  the  precipitate  has  collected  together.  Filter  and  wash  with 
water  containing  a  few  drops  of  nitric  acid.  Place  the  precipitate  in  a 
beaker,  put  several  pieces  of  granulated  zinc  on  it,  and  add  sufficient  dilute 
sulphuric  acid  to  cover  contents.  The  mixture  should  be  stirred  from 
time  to  time.  Equation? 

After  several  hours,  pour  off  the  acid,  remove  any  unchanged  zinc, 
wash  the  precipitate  with  water  by  decantation,  and  filter.  Dry  the 
brown  powder  between  sheets  of  filter  paper,  then  fuse  it  on  a  stick  of 
charcoal  by  directing  the  flame  of  a  blast-lamp  downward  upon  it  (?). 
Scour  the  metal  bead  with  sea  sand  ( ?). 

What  is  sterling  silver?  What  is  the  composition  of  standard  silver 
coin? 

Experiment  VII. — Precipitation  of  Silver  Sulphide. 

(a)  Pass  hydrogen  sulphide  through  2  or  3  cm.3  of  a  dilute  solution  of 
silver  nitrate  (?).     Equation?     Is  this  reaction  reversible  theoretically? 
To  a  portion  of  the  product  add  dilute  nitric  acid.     What  are  your  con- 
clusions now  as  to  the  reversibility  of  the  above  reaction  ? 

(b)  Expose  a  silver  coin  to  the  action  of  hydrogen  sulphide  (  ?).     What 
happens  chemically  when  silver  spoons  are  used  to  "beat"  an  egg? 
What  is  "oxidized  silver"? 

Silver  is  "tarnished"  by  the  sulphur  in  perspiration. 


SILVER.  3OI 

Experiment  VIII.— Action  of  Sodium  Thiosulphate  upon  Silver  Salts. 

Treat  small  quantities  of  each  of  the  following  substances  with  a  strong 
cold  solution  of  sodium  thiosulphate;  silver  chloride,  silver  bromide  and 
silver  sulphide  (?).  Equations?  Other  insoluble  salts  of  silver  behave 
in  a  similar  manner. 

Sodium  thiosulphate  ("hyposulphite  of  soda"  or  "hypo")  is  used 
largely  in  photograph^:  processes  for  fixing  the  image;  i.e.,  for  dissolving 
out  the  unreduced  silver  halide  on  the  photographic  plate. 

Experiment  IX. — Photography. 

Dip  a  piece  of  unsized  paper  (a  filter  paper  will  do)  in  a  sodium  chloride 
solution ;  allow  the  paper  to  drip  for  a  few  seconds,  then  dip  it  in  a  dilute 
solution  of  silver  nitrate.  Place  an  opaque  object  upon  the  paper  and 
expose  to  the  light  until  the  paper  around  the  object  has  a  deep  purple 
color,  then  dip  it  successively  into  a  solution  of  potassium-ferrous  oxalate* 
(a  mixture  of  solutions  of  ferrous  sulphate  and  potassium  oxalate),  dis- 
tilled water,  a  solution  of  sodium  thiosulphate  and  distilled  water.  Dry 
the  paper  and  fasten  it  in  your  laboratory  note-book.  The  foregoing 
procedure  is  similar  in  principle  to  that  involved  in  the  preparation  of  a 
"negative"  of  a  picture.  Explain  the  chemistry  of  each  step  in  the  pro- 
cedure. Equations  ? 

Photography  involves  two  processes — the  preparation  of  the  "  negative" 
and  the  printing  of  the  "positive."  The  negative  of  a  picture  is  usually 
taken  upon  a  prepared  plate  of  glass,  and  from  this  the  positive  is  printed 
upon  sensitized  paper.  The  negative  shows  the  "lights  and  shades" 
reversed,  hence  the  name.  The  positive  gives  the  objects  their  true 
appearance,  i.e.,  it  is  the  positive  of  the  original  object. 

The  modern  photographic  plate  usually  consists  of  a  glass  plate  covered 
with  a  gelatin  emulsion  in  which  is  suspended  finely  divided  silver 
bromide.  This  plate  is  exposed  in  a  camera  to  the  light  from  the  object 
to  be  photographed.  The  silver  salt  is  in  part  decomposed,  the  degree 
of  decomposition  being  proportional  to  the  intensity  of  the  light  and  the 
duration  of  the  exposure.  Thus  a  dark  object  will  reflect  few  light  rays, 
and  that  part  of  the  plate  exposed  to  these  rays  will  be  but  little  affected. 
A  white  surface  will  act  in  a  reverse  manner. 

According  to  one  view,  the  exposure  reduces  certain  portions  of  the 
silver  bromide  to  a  sub-bromide,  perhaps  Ag2Br,  as  indicated  by  the 
following  equations: 

2  Ag  Br  —  Ag2Br  +  Br. 

After  exposure,  there  is  usually  little  or  no  visible  alteration  in  the  film 
of  the  plate  as  the  decomposition  is  only  partial.  The  image  must  be 
developed.  This  is  done  by  washing  the  plate  with  a  reducing  agent, 

*A  sodium  carbonate  solution  of  pyrogallol  is  frequently  used  as  the  "developer" 
instead  of  the  potassium-ferrous  oxalate. 


302  EXPERIMENTAL  CHEMISTRY. 

such  as  an  alkaline  solution  of  pyrogallol  (photographer's  "pyro")  or  a 
solution  of  potassium-ferrous  oxalate.  This  solution,  which  is  known  as 
the  developer,  continues  the  action  initiated  by  the  light.  The  partly 
decomposed  silver  salt  is  affected  first,  and  with  a  speed  proportional  to  the 
intensity  of  the  illumination  undergone  by  each  part.  Some  of  the  salt 
is  reduced  to  metallic  silver: 


3Ag2Br  +  3FeC204  —  Fe2(C2O4)3  +  Fe«r3  +  Ag, 

The  depth  of  the  silver  deposit  is  thus  proportional  to  the  intensity  of 
the  light  upon  the  original  plate.  Those  portions  which  have  received 
the  most  light  have  the  larger  amount  of  the  silver  salts  decomposed  and 
are  dark  in  color. 

It  should  be  remembered  that  although  the  developer  acts  first  upon 
those  parts  reached  by  light,  it  can  reduce  the  whole  of  the  halide  upon 
the  plate.  When  the  relative  intensities  of  the  light  effects  —  i.e.,  the  "  con- 
trast" between  the  parts  variously  illuminated  —  have  been  brought  out 
sufficiently,  the  plate  is  removed  from  the  reducing  solution.  If  allowed 
to  remain  in  the  developer  long  enough,  all  the  silver  halide  would  be  re- 
duced, and  the  plate  would  be  uniformly  dark. 

The  plate  still  contains  that  portion  of  the  original  silver  salt  which  has 
not  been  acted  upon  by  the  light  or  the  developer.  This  must  be  re- 
moved before  the  plate  is  exposed  to  the  light;  otherwise,  it,  too,  will  be 
acted  upon  as  previously  described,  and  the  image  will  be  blurred  or 
obliterated  by  more  deposits  of  silver.  This  is  done  by  a  process  known 
as  fixing. 

When  it  is  seen  by  examination  that  the  development  has  proceeded 
long  enough,  the  plate  is  rinsed  in  water  and  placed  in  the  fixing  bath. 
This  is  a  solution  containing  sodium  thiosulphate  ("  sodium  hyposulphite  " 
or  "hypo"),  which  is  an  excellent  solvent  for  many  silver  compounds. 
The  fixing  bath  soon  dissolves  from  the  gelatin  film  the  silver  bromide 
which  remains  unaffected  by  the  light  or  the  developer.  The  "  chemistry  " 
of  the  "fixing"  process  may  be  represented  by  the  equation: 

2AgBr  +  3Na2S2O3  —  *  2NaAgS2O3,  Na2S2O3*  +  2NaBr. 

This  treatment  is  said  to  "clear"  or  "fix"  the  image.  The  plate  is 
no  longer  sensitive  to  light.  After  thorough  washing  it  is  allowed  to  dry. 
The  result  is  the  negative. 

The  plate  is  now  ready  to  be  used  in  making  prints.  This  is  done  by 
allowing  sunlight  to  fall  through  the  negative  upon  a  sheet  of  sensitized 
paper,  which  may  be  regarded,  for  purposes  of  discussion,  as  paper 
covered  with  a  film  of  albumen  holding  a  deposit  of  silver  chloride.  The 
negative  and  the  paper  are  placed  film  to  film.  The  action  of  sunlight 
upon  this  paper  is  similar  to  that  upon  the  plate.  Exposure  liberates 
silver. 

2AgCl->Ag2Cl+_Cl 
AgCl->Ag+Cl. 

*  Formula  of  the  crystals  of  the  complex  salt  obtained  from  the  solution. 


SILVER.  303 

As  the  dark  portions  of  the  negative  which  correspond  to  the  light 
parts  of  the  object  transmit  the  fewer  light  rays  (sunlight  cannot  penetrate 
the  parts  over  which  there  is  a  deposit  of  silver),  it  is  obvious  that  what  is 
dark  in  the  negative  will  be  light  in  the  positive  and  vice  versa.  The 
print  will  have  the  same  shading  as  the  object. 

The  print  is  then  toned  and  fixed.  Toning  consists  in  imparting  to  the 
print  a  rich  color  by  replacing  part  of  the  Ag3Cl  with  gold  from  a  solu- 
tion of  sodium  chlor-aurate,  NaAuCl4. 

3Ag2Cl  +  NaAuCl4  —  NaCl  +  6AgCl  +  Au. 
or 

2Ag2Cl  +  K2PtCl4— »  2KC1   +  4AgCl  +  Pt. 

Fixing  removes  the  unchanged  silver  salt,  the  operation  being  per- 
formed as  previously  described.  After  thorough  washing  the  picture  is 
dried  and  mounted. 

"Retouching"  is  a  process  whereby  blemishes  are  removed  from  both 
the  negative  and  the  positive  by  the  use  of  India  ink  or  colored  pencils. 

Experiment  X. — Silver  Nitrate  and  Organic  Matter. 

(a)  Press  a  small  crystal  of  silver  nitrate  between  the  thumb  and  fore- 
finger for  icor  15  seconds.     After  several  hours,  examine  your  finger  (?). 

(b)  Using  a  solution  of  silver  nitrate,  write  your  name  upon  a  sheet  of 
paper  with  a  pointed  instrument.     Expose  the  writing  to  the  direct  sun- 
light  (?). 

What  is  lunar  caustic  ?     What  is  the  origin  of  this  name  ? 

Experiment  XI. — Precipitation  of  Silver  Chr ornate. 

Add  a  solution  of  potassium  dichromate  to  a  silver  nitrate  solution. 
Describe  the  result.  Equation? 

Experiment  XII. — Interaction  Silver  Salts  and  Potassium  Cyanide. 

(a)  To  a  solution  of  silver  nitrate  add  slowly,  drop  by  drop,  a  potassium 
cyanide  solution  (poison)  (?).     Equation?     Add  the  precipitant  in  ex- 
cess   (?).     Pass    hydrogen    sulphide    through    this    solution.     Result? 
Is  the  silver  ion  present  ?     Explain.     Equation  ?. 

What  conclusion  with  regard  to  the  formation  of  complex  ions  may  be 
generally  drawn  when  the  precipitate  dissolves  in  excess  of  the  precipi- 
tant? Give  examples. 

(b)  All  silver  salts  are  soluble  in  potassium  cyanide.     Test  (?). 

The  bath  ordinarily  used  for  silver  electroplating  consists  of  a  solution 
of  potassium  argenticyanide.  Solutions  of  the  simple  salts  do  not  give 
a  coherent  film. 

AgNO3  +  2KCN    <=*  KAg(CN)2  +  KNO3 
KAg(CN)2    <=>  K-  +  Ag(CN)'2 
Ag(CN)'2      <=>Ag- 
Ag'  +  e       <=±  Ag. 


304  EXPERIMENTAL  CHEMISTRY. 

Experiment  XIII. — Analytical  Reactions  of  Silver  Salts. 

(a)  Compounds  of  silver  heated  on  charcoal  before  the  blow-pipe  give 
a  white  metallic  globule  (?).     Test  (?). 

(b)  Dissolve  a  small  crystal  of  chemically  pure  silver  nitrate  in  dis- 
tilled water.     Test  the  reaction  of  the  solution  with  litmus  paper  ( ?).     Is 
the  salt  hydrolyzed  by  water  ?     In  this  respect  it  approaches  what  group 
of  metals? 

What  are  your  conclusions  as  to  the  relative  chemical  activity  of  silver 
and  copper  as  metallic  elements? 

(c)  Those  reactions  which  are  of  particular  importance  in  analytical 
chemistry  are  given  in  Exps.  V  and  XII. 

(d)  Silver  chloride  is  insoluble  in  either  hot  or  cold  water. 

GOLD,    All. 

At.  Wt.   197.2  Sp.   Gr.    19.3. 

Experiment  I. — Properties  of  Gold. 

(a)  Physical  properties.     Examine  a  piece  of  gold  and  state  briefly 
the  most  obvious  physical  properties  of  the  metal.     (The  student  is  re- 
minded that  the  "commercial"  metal  is  an  alloy  of  gold  and  copper. 
The  latter  gives  it  greater  hardness). 

Record  its  melting  point,  specific  heat  and  atomic  heat.  Calculate 
the  atomic  weight  of  gold  by  Dulong  and  Petit's  Law.  Pure  gold  is 
"24-carat."  Explain  the  use  of  the  word  "carat."  What  is  the 
composition  of  the  American  standard  gold  coin? 

(b)  Chemical   properties.     Is   gold   affected   by   free   oxygen  ?     Give 
reasons  for  your  answer. 

What  is  the  valency  of  gold? 

Direct  a  stream  of  hydrogen  sulphide  against  a  piece  of  pure  gold  ( ?). 
The  negative  result  is  indicative  of  its  inacitivity.  Gold  is  the  least  active 
of  all  of  the  familiar  metals. 

Does  gold  displace  hydrogen  from  dilute  acids?  Before  attempting  to 
answer,  see  table  of  "Solution  Tensions." 

Gold  combines  with  free  chlorine,  therefore,  its  solubility  in  aqua 
regia  with  the  formation  of  auric  chloride,  AuCl3. 

Gold  does  not  interact  with  any  of  the  oxacids  except  selenic  acid. 

Experiment  II. — Analytical  Reactions  of  Salts  of  Gold. 

(a)  All  compounds  of  gold  are  decomposed  by  heat  with  liberation  of 
the  metal.     Test,  using  a  solution  of  auric  chloride  (?). 

(b)  Hydrogen  sulphide  with  solutions  of  AuCl3  gives  a  black  precipi- 
tate, auro-auric  sulphide  (Au2S,  Au2S3).      This  sulphide  is  insoluble  in 
either  hydrochloric  or  nitric  acid,  but  soluble  in  aqua  regia.     It  is  also 
soluble  in  ammonium  sulphide,  forming  an  ammonium  sulpho-  or  thio- 
aurate. 

(c)  Ferrous  sulphate  gives  a  brown  or  purple  precipitate  of  metallic 


GOLD.  305 

gold.     If  the  solution  of  the  auric  salt  is  dilute  a  bluish  color  is  imparted 
to  the  mixture. 

AuCl3  +  3FeSO4->Fe2(SOJ3  +  FeCl3  +  Au. 

(d)  Stannous  chloride  gives  with  solutions  of  auric  chloride  a  precipi- 
tate or  coloration  which  varies  in  color  from  reddish-brown  to  purple. 
The  color  is  due  to  the  presence  of  a  compound  of  uncertain  composition, 
known  as  the  purple  oj  Cassius. 

The  test  may  be  made  as  follows:  Add  a  ferric  chloride  solution  to  a 
solution  of  stannous  chloride  until  the  latter  has  a  permanent  yellow  color. 
Add  a  drop  of  this  solution  to  an  auric  chloride  solution  ( ?),  or  dip  a  glass 
rod  into  the  former  solution,  and  then  into  the  latter  (?). 

(e)  Other  reducing  agents,  like  oxalic  acid  and  potassium  nitrite,  when 
warmed  with  solutions  of  AuCl3,  precipitate  the  gold.     If  the  former  re- 
agent is  used,  the  gold  is  deposited  in  a  spongy  form  similar  to  that  used 
by  dentists. 

2AuCl3  +  3H2C2O4  ->  6HC1  +  6CO2  +  2Au. 

2AuCl3  +  3KNO2  +  3H2O  —  3KNO3  +  6HC1  +  2Au. 

PROBLEMS. 

1.  How  much  cupric  oxide  is  formed  by  heating  10  grams  of  copper  in 
the  air? 

2.  How  much  cuprous  oxide  can  be  prepared,  theoretically,  by  boiling 
10  grams  of  cupric  sulphate  with  a  solution  of  potassium  hydroxide  and 
glucose  ? 

2CuSO4.sH2O  +  4KOH-*  Cu2O  +  O  +  2K2SO4  +  i2H2O. 

3.  Calculate  the  percentage  of  copper  in  malachite,  Cu(OH)2,Cu(CO3) 
In  azurite,  Cu(OH)2,2CuCO3. 

4.  If  2  grams  of  silver  chloride  yield  1.505  grams  of  silver,  what  is  the 
atomic  weight  of  silver? 

5.  Calculate  the  quantities  of  silver  and  gold  which  can  be  deposited 
from  their  respective  solutions  by  the  gram-atomic  weight  of  copper.     By 
i  gram  equivalent.     By  i  gram  of  copper. 

6.  The  same  electric  current  passes  through  separate  solutions  of  silver 
nitrate  and  copper  sulphate.      The  electrodes  in  each  case  are  plates  of 
platinum.     The  cathode  in  the  copper  sulphate  solution  increased  0.636 
gram  in  weight.     How  much  did  the  cathode  in  the  silver  nitrate  solution 
increase  in  weight? 

7.  One  coulomb  of  electricity  deposits  0.0011175  gram  of  silver.     This 
quantity  is  known  as  the  electro-chemical  equivalent  of  silver.     How  many 
coulombs  are  required  for  the  deposition  of  one  gram  equivalent  of  silver  ? 
Of  any  substance  ?     Calculate  the  electro-chemical  equivalent  of  copper. 


CHAPTER  XXVI. 
ALKALINE  EARTH  METALS. 

(Glucinum,  Gl.  9.1) 
Magnesium,  Mg.  24.36 
Calcium,  Ca.  40 .  i 

Strontium,  Sr.         87  . 6 
Barium,  Ba.  137  .4 

(Radium,  Rd.       225.) 

The  metals  of  this  family  are  called  the  "  alkaline-earth"  metals  because 
they  form  the  transition  from  the  "alkalies"  to  the  "earth"  metals,  such 
as  aluminum,  gallium,  etc. 

None  of  the  elements  of  the  family  occur  in  nature  in  the  uncombined 
condition,  and  have  only  in  more  recent  years  been  prepared  from  their 
compounds.  With  the  exception  of  magnesium,  they  are  seldom  isolated. 
In  the  case  of  glucinum  this  is  due  largely  to  the  comparative  rarity  of 
its  compounds.  Calcium,  strantium  and  barium  whose  compounds 
are  abundant,  are  seldom  prepared  in  the  pure  state  save  in  very  small 
quantities  because  of  the  extreme  difficulty  in  isolating  them,  and  owing 
also  to  the  fact  that  they  are  of  little  commercial  importance. 

Although  these  elements  are  brought  into  the  same  family  by  the  peri- 
odic grouping  and  possess  many  marked  similarities,  yet  they  do  not  form 
an  altogether  homogeneous  and  coherent  group. 

Glucinum  and  magnesium  are  white  metals  which  do  not  rust  (oxidize) 
rapidly  in  the  air,  while  calcium,  strontium  and  barium,  under  similar 
conditions,  quickly  lose  their  silver-white  luster.  Glucinum  does  not 
decompose  water;  magnesium  displaces  hydrogen  from  boiling  wrater; 
and  the  other  metals  decompose  water  spontaneously  at  ordinary  temper- 
atures, forming  hydroxides. 

Glucinum  and  magnesium  resemble  zinc  and  cadmium  and  differ  from 
calcium,  strontium  and  barium  in  that  their  sulphates  are  readily  soluble 
in  water,  but  these  relations  are  reversed  in  the  case  of  their  sulphides 
which  are  hydrolyzed  by  water.  Glucinum  differs  from  magnesium  and 
resembles  zinc  in  that  its  hydroxide  is  acidic  as  well  as  basic,  i.e.,  is  soluble 
in  sodium  or  potassium  hydroxide. 

Calcium,  strontium  and  barium  bear  a  much  closer  resemblance  to 
each  other  in  most  of  their  physical  and  chemical  properties  than  to  either 
glucinum  or  magnesium.  Their  hydroxides  show  an  increasing  solubility* 
and  alkalinity  in  the  order  named.  These  solutions,  which  are  very 
dilute,  owe  their  strong  alkalinity  to  the  high  degree  of  ionization  (see 

*  200,  630,  and  2,200  parts  of  the  hydroxides,  respectively,  dissolve  in  1,000,000 
parts  of  water. — Smith. 

306 


ALKALINE    EARTH    METALS.  307 

table)  which  they  undergo.  The  order  of  the  solubility  of  the  sulphates 
of  these  three  metals  is  reversed — 2100,  no  and  2.3  parts,  respectively, 
being  soluble  in  1,000,000  parts  of  water. 

All  of  these  metals  form  an  oxide  of  the  type  MO;  calcium,  strontium 
and  barium  also  form  peroxides  of  the  type  MO2.  The  oxides  combine 
with  water,  forming  hydroxides. 

When  the  nitrates  of  calcium,  strontium,  and  barium  are  heated  they 
are  decomposed,  yielding  the  oxides  of  the  metals,  nitrogen  peroxide  and 
oxygen. 

The  carbonates  which  are  insoluble,  when  heated,  are  decomposed  into 
the  oxide  of  the  metal  and  carbon  dioxide. 

The  metals  of  the  family  usually  exhibit  a  valence  of  two. 

Radium,  one  of  the  recently  discovered  elements,  has  never  been  iso- 
lated. It  is  a  "radio-active"  substance  which  appears  to  belong  to  this 
family.  It  exhibits  a  valence  of  two  in  those  of  its  salts  which  have  been 
investigated. 

GLUCINUM,    Gl. 

At.  Wt.  9.1     Sp.  Gr.  1.8. 
See  lecture-notes  or  text-book. 

MAGNESIUM,    Mg. 

At.  Wt.  24.36     Sp.  Gr.  1.74. 
Experiment  I. — Properties  of  Magnesium. 

(a)  Examine  the  elementary  substance  in  the  forms  of  "ribbon, "  wire 
and  powder.     Scrape  a  piece  of  magnesium  ribbon,  and  note  its  color  and 
luster. 

(b)  Combustion  of  magnesium  in  the  air.     Introduce  a  piece  of  the 
ribbon  into  the  flame  with  the  forceps  ( ?).     What  is  the  product  ?     Equa- 
tion? 

Apply  the  flame  (caution)  to  a  small  pinch  of  magnesium  powder  upon 
an  iron  plate  ( ?).  Account  for  the  difference  in  the  speeds  of  the  reactions. 

Note. — Powdered  magnesium  is  frequently  used  as  one  of  the  com- 
ponents of  "flash-light  powder"  for  use  in  photography. 

(c)  Solubility  in  acids.     Treat  individually  small  pieces  of  magnesium 
ribbon  with  dilute  sulphuric  acid  ( ?),  hydrochloric  acid  ( ?)>  nitric  acid  ( ?) 
and  acetic  acid  ( ?).     Express  each  reaction  by  an  equation. 

Experiment  II. — Magnesium  Salts. 

Examine  the  different  salts  of  magnesium  (end  shelf),  noting  the  ob- 
vious physical  properties.  Give  the  name,  formula,  color  and  solubility* 
of  each  salt.  Tabulate  the  foregoing  data. 

*See  Comey's  Dictionary  of  Chemical  Solubilities,  and  SeidelPs  Solubilities  of 
Inorganic  and  Organic  Substances. 


308  EXPERIMENTAL  CHEMISTRY. 

Experiment  III. — Hydrolysis  of  Magnesium  Chloride. 

Ascertain  by  experiment  whether  magnesium  chloride  dissolves  com- 
pletely in  water  (?).  Test  the  solution  with  litmus  paper  (?).  Explain. 
Equation  ? 

Place  several  small  crystals  of  the  salt  in  a  dry  test  tube,  and  heat 
strongly.  Test  the  reaction  of  the  evolved  vapors,  and  the  liquid  which 
condenses  in  the  tube  toward  litmus  paper  (?).  Explain.  Equation? 

Give  the  formula  for  carnallite. 

Experiment  IV. — Precipitation  of  Magnesium  Hydroxide  and  its  Solu- 
bility in  Ammonium  Compounds. 

(a)  To  10  cm.3  of  solution  of  magnesium  sulphate  add  a  solution  of 
sodium  hydroxide  or  potassium  hydroxide  (?).     Equation? 

(b)  Repeat  (a)  using  ammonium  hydroxide  as  the  precipitant.     Ob- 
serve the  character  of  the  precipitate.     Express  the  reaction  by  an  equa- 
tion.    Is  the  precipitation  of  the  magnesium  complete  ?     Before  attempt- 
ing to  answer  this  last  question,  add  a  strong  solution  of  ammonium 
chloride  (note  that  this  is  one  of  the  "products"  in  the  above  reaction) 
in  large  excess.     Explain  the  disappearance  of  the  precipitate  in  terms 
of  the  ion-  or  solubility-product  constant. 

(c)  Can  magnesium  hydroxide  be  precipitated  in  the  presence  of  an  ex- 
cess of  an  ammonium  salt? 

Experiment  V. — Precipitation  of  Hydrated  Basic  Carbonate  of  Mag- 
nesium; Its  Solubility  in  Ammonium  Compounds. 

(a)  Add  a  slight  excess  of  a  sodium  carbonate  solution  to  10  cm.3  of 
a  solution  of  magnesium  sulphate.     The  composition  of  the  precipitate 
varies  with  conditions. 

(b)  Repeat  (a),  using  ammonium  carbonate  instead  of  sodium  carbonate. 
Add  an  excess  of  a  strong  solution  of  ammonium  chloride.     Result? 
Equation  ? 

(c)  Can  the  basic  carbonate  of  magnesium  be  precipitated  in  the  pres- 
ence of  an  excess  of  an  ammonium  salt  ? 

Experiment  VI. — Preparation  of  Magnesium  Chloride. 

To  15  cm.3  of  a  dilute  solution  of  hydrochloric  acid  add  an  excess  of 
basic  magnesium  carbonate,  filter,  evaporate  to  small  volume  and  crystal- 
lize. Equation  ?  What  is  the  formula  of  the  crystallized  salt  ? 

Experiment  VII. — Analytical  Reactions  of  Magnesium  Salts.  Precipi- 
tation of  Ammonium  Magnesium  Phosphate. 

To  10  cm.3  of  a  solution  of  a  soluble  magnesium  salt  add  an  excess  of 
ammonium  chloride,  then  add  cautiously  2  cm.3  or  3  cm.3  of  ammonium 
hydroxide  or  an  equal  volume  of  a  solution  of  ammonium  carbonate. 
(Why  do  neither  of  the  latter  substances  produce  a  precipitate?)  To 
the  clear  solution  thus  obtained  add  a  disodium  hydrogen  phosphate 


ALKALINE  EARTH  METALS.  309 

(Na2HPO4)  solution  until  precipitation  ceases.     Describe  the  precipitate. 
Express  the  reaction  by  an  equation. 

When  ammonium  magnesium  phosphate  is  ignited  it  is  converted  into 
the  anhydrous  pyrophosphate  of  magnesium. 

(2NH4MgP04,  6H20  ->  Mg2P207  +  2NH3  +  13^0). 
What  is  the  color  and  valence  of  the  magnesium  ion  ? 

CALCIUM,   ca. 

At.  Wt.  40.1       Sp.  Gr.  1.85. 
Experiment  I. — Properties  of  Calcium. 

(a)  Enumerate  the  most  obvious  physical  properties  of  the  metal. 

(b)  Does  it  interact  with  water?     Equation? 

(c)  Will  it  burn  in  the  air?     Equation? 

Note. — The  product  of  the  combustion  is  a  mixture  of  the  oxide  and 
the  nitrite  (Ca3N2).  Compare  with  the  combustion  of  magnesium  in  air. 

Experiment  II. — Action  of  Acids  upon  Calcium  Carbonate. 

(a)  Recall  or  repeat  Exp.  II  (c),  "Carbon."     Indicate  by  an  equation 
the  action  of  hydrochloric  acid  upon  marble  (impure  calcium  carbonate). 

(b)  Test  a  small  quantity  of  powdered  marble  with  dilute  sulphuric 
acid,  and  note  the  effect.     Equation  ? 

(c)  Repeat  (b)  substituting  nitric  acid  for  sulphuric  acid. 

Experiment  III. — Preparation  of  Calcium  Oxide  (Quicklime). 

Ignite  about  3  grams  of  powdered  marble  for  15  to  20  minutes  in  an 
open  crucible  placed  upon  a  pipe-stem  triangle.  Maintain  the  crucible 
and  contents  at  the  highest  temperature  obtainable  with  the  aid  of  a  blast- 
lamp.  Occasionally  stir  the  mass  with  a  platinum  wire.  When  the  mass 
has  cooled,  add  about  10  cm.3  of  water;  stir  the  mixture  (?).  Test  reac- 
tion of  the  liquid  with  litmus  paper  (?).  Does  water  act  upon  marble? 
Add  a  few  drops  of  hydrochloric  acid  to  the  above  mixture  (?).  How 
do  the  results  compare  with  the  action  of  acid  upon  marble  ?  What  are 
your  conclusions  as  to  the  identity  of  the  substance  in  the  crucible  ?  The 
decomposition  of  the  marble  by  heat  is  a  reversible  reaction.  How  may 
it  be  brought  to  a  condition  of  equilibrium?  How  may  it  be  reversed? 
How  may  it  be  reversed  without  altering  the  temperature?  What  pre- 
cautions must  be  taken  in  the  construction  of  a  lime-kiln  that  the  complete 
decomposition  of  marble  into  calcium  oxide  and  carbon  dioxide  may  be 
insured?  Indicate  by  equations  the  chemical  equilibria  involved  in  the 
above  decomposition. 

Experiment  IV. — Preparation  and  Properties  of  Calcium  Hydroxide. 
Lime-Water. 

(a)  Pour  5  cm.3  of  distilled  water  upon  8  grams  of  quicklime  in  a  porce- 


310  EXPERIMENTAL  CHEMISTRY. 

lain  dish,  and  note  the  results  (?).  The  foregoing  process  is  known  as 
the  "slaking"  or  "slacking"  of  lime.  The  product  is  known  as  calcium 
hydroxide.  Express  the  reaction  by  an  equation.  What  is  "mortar"? 
What  is  "air-slacked"  lime? 

(b)  Solubility*  and  effect  of   calcium  hydroxide  upon  litmus.     Put 
4  or  5  grams  of  the  calcium  hydroxide  prepared  in  (a)  into  a  flask  or 
bottle  containing  200  cm. 3  of  distilled  water,  cork  tightly,  shake  vigorously 
from  time  to  time,  and  allow  the  mixture  to  stand  until  the  solution  has 
become  clear,  when  the  latter  may  be  decanted,  drawn  off  with  a  siphon  or 
removed  with  a  pipette.     If  too  much  time  is  consumed  in  allowing  the 
mixture  to  settle,  it  may  be  filtered  rapidly,  and  the  clear  liquid  (lime- 
water)  used  for  the  following  experiments. 

(c)  Test  the  action  of  the  solution  upon  litmus  paper  (?).     The  result 
indicates  the  presence  of  what  ion  ?     Explain. 

(d)  Recall  or  repeat  Exp.  II  (d),  "Carbon."     Equations? 

To  a  piece  of  old  mortar  in  a  test  tube  add  dilute  hydrochloric  acid. 
Identify  the  gas  (?).  What  chemical  action  is  involved  in  the  "harden- 
ing" of  mortar? 

(e)  Recall  or  repeat  Exp.  V  "Carbon."     Equations?     When  is  water 
said  to  possess  "  temporary  hardness"  ?     It  is  due  to  the  presence  of  what 
salt? 

Experiment  V. — Calcium  Salts. 

Repeat  Exp.  II  "Magnesium,"  substituting  the  word  calcium  where 
the  word  magnesium  appears. 

Experiment  VI. — Preparation  of  Calcium  Salts. 

(a)  Calcium  chloride.     Suggest  by  an   equation   such   a  method  for 
the  preparation  of  calcium  chloride  from  one  of  its  salts  as  will  not  in- 
volve the  simultaneous  formation  of  any  other  salt,  acid  or  base.     Is 
calcium  chloride  deliquescent  or  efflorescent?     It  is  frequently  used  for 
what  purpose  in  connection  with  the  manipulation  of  gases? 

(b)  Calcium  hypochlorite.     Describe  a  method  for  the  preparation  of 
this  salt.     (Suggestion,  see  Exp.  Ill  (a),  "  Chlorine.")    Equation  ?    What 
is  "bleaching  powder "?f     Its  formula? 

Experiment  VII. — Dehydration  of  Gypsum.     Plaster  of  Paris. 

(a)  Heat  a  few  grams  of  powdered  gypsum  or  a  small  piece  of  a  selenite 
crystal  (CaSO4.  2H2O)  in  a  test  tube  and  record  the  obvious  results. 
Pulverize  the  residue  and  ascertain  whether  it  will  become  solid  when 

*  At  1 8°,  600  parts  of  water  by  weight  dissolve  i  part  of  the  hydroxide;  at  100°, 
about  twice  as  much  water  is  required  to  dissolve  the  same  quantity  of  salt. — Seidtll. 

fThe  probable  formula  of  bleaching  powder  is  CaCl  (CIO)— a  mixed  salt  rather 
than  an  equi-molar  mixture  of  CaCl2  andCa(OCl)2.  This  view  is  supported  by  the 
fact  that  CaCl2  cannot  be  dissolved  out  with  alcohol,  nor  is  the  salt  deliquescent  as  is 
calcium  chloride  (CaCl2).  The  ions  Ca",  Cl',  and  CIO'  are  present  in  a  solution  of 
the  salt. 


ALKALINE  EARTH  METALS.  311 

mixed  with  a  little  water  to  form  a  paste  and  allowed  to  stand  (?).     See 
whether  gypsum  itself  will  act  in  the  same  way  with  water  ( ?) .     Equations  ? 

Note. — If  all  of  the  water  is  removed  from  gypsum  by  heating  or  the 
temperature  is  allowed  to  rise  much  above  125°,  the  product  when  mixed 
with  water  does  not  "set"  quickly. 

(b)  Make  a  thick  paste  by  mixing  a  little  water  with  "plaster  of  Paris." 
Place  the  paste  upon  a  glass  plate  and  make  a  cast  of  a  coin.  Chemically, 
what  is  "plaster  of  Paris?"  Probable  formula? 

Experiment  VIII. — Solubility  of  Phosphates  of  Calcium. 

(a)  Ascertain  whether  or  not  the  tertiary  orthophosphate  of  calcium 
(Ca3(PO4)2)    is    soluble    in    water?     Would  such  a  salt  make  a  good 
fertilizer  ?     Reasons  for  your  answer  ? 

(b)  Repeat     (a)  using    primary    calcium     phosphate     (Ca(H2PO4)2. 
What  is  "superphosphate  of  lime"?     Its  use? 

Experiment  IX. — Analytical  Reactions  of  Calcium  Salts. 

(a)  Flame  test.     Introduce  a  small  quantity  of  calcium  chloride  into 
the  Bunsen  flame  by  means  of  a  platinum  wire  (?).     Examine  the  flame 
with  a  spectroscope,  noting  particularly  the  presence  of  two  bands — a 
red  and  a  green  one — which  impart  the  brick-red  color  to  the  flame. 

(b)  (Use  a  dilute  solution  of  pure  calcium  chloride  for  the  following 
tests.)     Examine  such  solutions  of  calcium  salts  as  are  found  on  the  end 
shelf.     What  is  the  color  of  the  calcion?     Its  valence?     Justify  your 
conclusions. 

(c)  Calcium  carbonate.     To  a  portion  of  the  solution  containing  the 
calcion  add  a  slight  excess  of  an  ammonium  carbonate  solution  (?). 
Warm  the  solution  if  necessary  to  procure  complete  precipitation.     Ex- 
press the  reaction  by  an  equation.     Filter,  spread  the  filter  paper  with 
precipitate  upon  a  glass  plate  and  divide  the  precipitate  into  two  parts. 
Treat   one   part  with   hydrochloric   acid    ( ?)    and  the  other  with  dilute 
acetic  acid  (?).     Equations?     What  are  your  conclusions  as  to  the  rela- 
tive strength  of  acetic  and  carbonic  acids?     Confirm  your  inference  by 
giving  the  degrees  of  ionization  in  .1  normal  solutions  of  the  acids.     Ex- 
plain in  terms  of  the  "electrolytic  dissociation  theory"  how  the  insoluble 
salt  of  a  "weak"  acid  is  dissolved  by  a  stronger  acid. 

(d)  Calcium  oxalate.     Add  an  excess  of  an  ammonium  oxalate  solu- 
tion to  a  dilute  solution  of  calcium  chloride  (?).     Equation?    Filter  and 
divide  precipitate  into  two  parts.     Test  the  precipitate  as  in  (c).     Inter- 
pret results. 

Explain  the  difference  in  behavoir  of  the  oxalate  and  carbonate  of 
calcium  toward  acetic  acid,  taking  into  account  the  difference  in  solubility 
of  the  two  salts,  and  the  behavior  of  oxalic  acid  and  carbonic  acid,  re- 
spectively. Why  was  ammonium  oxalate  used  as  the  preciptant  in  pref- 
erence to  oxalic  acid  (recall  or  repeat  Exp.  II  (c~),  "Chemical  Equilib- 


312  EXPERIMENTAL  CHEMISTRY. 

rium")?  Under  what  conditions  may  calcium  oxalate  be  precipitated 
from  a  calcium  chloride  solution  ?  Explain. 

(e)  I.  Calcium  sulphate.  (e,I.  Repeat  Exp.  II  (ft),  "Chemical  Equi- 
librium.") 

(e)  II.  To  a  dilute  solution  of  calcium  chloride  add  a  slight  excess 
of  dilute  sulphuric  acid  ( ?)  Equation  ?  Filter  and  divide  the  nitrate 
into  two  parts.  Add  sufficient  ammonium  hydroxide  to  one  portion  to 
barely  neutralize  the  free  acids  (determine  the  "end  point"  by  means 
of  litmus  paper).  Now  add  ammonium  oxalate  to  the  neutralized  solu- 
tion (?).  Explain  the  precipitation  of  calcium  oxalate.  Which  is  the 
more  soluble — the  sulphate  or  oxalate  of  calcium  ?  Give  the  actual  solu- 
bilities of  each.  (See  works  to  which  you  have  been  referred  previously.) 

(e)  III.  "Permanent  hardness"  of  water.  Boil  the  second  portion  of 
the  nitrate  from  (e),  II.  Compare  the  negative  result  with  that  obtained 
in  Exp.  IV  (e).  To  the  above  solution  add  a  solution  of  sodium  carbo- 
nate (?).  Equation?  Devise  a  method  for  proving  that  the  precipitate 
is  calcium  carbonate  (?).  Is  the  sulphate  or  the  carbonate  more  solu- 
ble (?).  Reasons  for  your  answer?  Give  the  actual  solubilities  of  the 
two  substances.  Can  "permanent  hardness"  be  removed  by  boiling? 
Give  one  method  for  removing  it.  What  are  some  of  the  objections  to 
"hard  water"  for  domestic  and  technical  use? 


STRONTIUM,    Sr. 

At.  Wt.  87.6          Sp.   Gr.   2.54. 
Experiment  I. — Properties  of  Strontium  and  its  Salts. 

Strontium  is  a  yellowish-white,  rather  tough  metal  which  slowly  oxi- 
dizes spontaneously  in  the  air.  It  burns  when  heated  in  the  air,  and  de- 
composes water  at  ordinary  temperatures.  It  is  malleable  and  ductile  and 
fuses  at  red  heat. 

The  physical  and  chemical  properties  of  the  compounds  of  strontium  are 
very  similar  to  those  of  calcium. 

Experiment  II. — Strontium  Salts. 

Repeat  Exp.  II,  "Magnesium,"  substituting  the  word  strontium  for 
magnesium. 

Experiment  III. — Analytical  Reactions  of  Strontium  Salts. 

(a)  Flame  test.     Dip  a  platinum  loop  into  a  solution  of  strontium 
chloride  and  hold  it  in  the  Bunsen  flame  ( ?).     Examine  the  flame  with  a 
spectroscope  and  make  a  sketch  of  the  spectrum  showing  the  position  of 
the  lines  with  reference  to  the  sodium  lines. 

Note. — Anhydrous  strontium  is  mixed  with  charcoal,  sulphur  and  potas- 
sium chlorate  to  make  "red  fire." 

(b)  Strontium  carbonate.     Add  ammonium  hydroxide  and  ammonium 


ALKALINE    EARTH    METALS.  313 

carbonate  to  a  dilute  solution  of  strontium  chloride  (?).  Equation? 
Filter,  wash  the  precipitate  and  test  the  action  of  acetic  and  hydrochloric 
acids  upon  it  (?).  Equations? 

(c)  Strontium  sulphate.  To  a  solution  containing  strontium  add  a 
small  amount  of  a  solution  of  calcium  sulphate;  heat  to  boiling,  and  if  no 
precipitate  appears  immediately,  let  the  mixture  stand  for  10  or  15  minutes. 
Explain  why  the  precipitate  forms  so  slowly.  What  do  you  infer  regard- 
ing the  relative  solubilities  of  calcium  and  strontium  sulphates.  Give 
the  actual  solubilities  of  each. 

What  is  the  color  of  the  strontium  ion  ?     Its  valence  ? 

BARIUM,    Ba. 

At.  Wt.   137.4         Sp.   Gr.  3.6. 
Experiment  I. — Properties  of  Barium  and  its  Salts. 

Barium  is  a  bright  yellow  metal  which  readily  oxidizes  spontaneously 
in  the  air.  It  decomposes  water  energetically  at  ordinary  temperatures. 
The  metal  is  malleable  and  ductile  and  fuses  at  red  heat. 

The  physical  and  chemical  properties  of  the  compounds  of  barium 
resemble  those  of  calcium  and  strontium. 

Experiment  II. — Barium  Salts. 

Same  as  Exp.  II,  "Strontium.  " 

Experiment  III. — Water  of  Hydration  in  Barium  Chloride. 

A  crucible  with  the  lid,  which  has  been  cleaned,  is  placed  on  a  pipe- 
stem  triangle  and  heated  with  the  Bunsen  burner  for  a  few  minutes. 
Cool,  and  determine  carefully  the  weight  of  the  crucible.  Weigh  into 
crucible  accurately  about  3  grams  of  pure  crystallized  barium  chloride. 
Place  the  crucible  with  its  contents  in  an  air  bath  at  120°  to  130°  C.  for 
about  one  hour,  then  cool  for  10  minutes  and  weigh;  or  the  crucible  and  its 
contents  may  be  placed  on  a  triangle  and  heated  gently  with  a  small  Bunsen 
burner  flame.  The  temperature  is  gradually  raised  until  the  crucible 
attains  a  low  red  heat  at  which  it  is  maintained  for  10  minutes  when  it  is 
cooled  and  weighed  as  per  above  directions.  The  heating  and  weighing 
is  repeated  until  the  weight  is  constant. 

Record  your  results  as  follows:  Grams. 

Weight  of  Xble  +  barium  chloride  = 
Weight  of  Xble  alone  = 

Weight  of  barium  chloride  taken  = 
Weight  of  Xble  +  contents  before  heating  = 
Weight  of  Xble  +  contents  after  heating  = 

Weight  of  water  found  = 
Per  cent,  of  water  in  barium  chloride  = 
Theoretical  per  cent.  = 

Error  = 


314  EXPERIMENTAL  CHEMISTRY. 

Experiment  IV. — Analytical  Reactions  of  Barium  Salts. 

(a)  Flame  test.     Same  as  Exp.  Ill  (a),  "Strontium." 

Note: — Barium  chlorate  is  used  with  charcoal  and  sulphur  to  prepare 
"green  fire." 

(b)  Barium  carbonate.     Same  as  Exp.  Ill,  b,  "Strontium." 

(c)  Barium  chromate.     To  a  dilute  solution  of  barium  chloride  add  a 
solution  of  potassium  chromate  (?).     Ascertain  whether  the  precipitate 
is  soluble  in  dilute  acetic  acid.     In  hydrochloric  acid. 

Note. — Neither  calcium  nor  strontium  chromate  is  precipitated  from 
dilute  solutions  acidified  with  acetic  acid. 

(d)  Barium  sulphate.     Add  dilute  sulphuric  acid  to  a  barium  chloride 
solution  (?).     Equation? 

To  10  cm. 3  of  a  solution  of  barium  chloride  add  20  cm.3  of  a  clear 
saturated  solution  of  strontium  sulphate  (prepared  by  shaking  the  salt 
with  distilled  water  and  filtering)  (?).  Which  is  the  more  soluble — stron- 
tium or  barium  sulphate? 

Arrange  the  sulphates  of  calcium,  strontium  and  barium  in  the  order 
of  decreasing  solubility.  Give  the  actual  solubility  of  each. 

Experiment  V. — Detection  of  the  Alkaline  Earth  Metals  in  a  Mixture. 

(a)  Give  two  methods  for  distinguishing  the  compounds  of  the  metals 
of  this  group. 

(b)  How  may  the  carbonate  of  magnesium  be  separated  from  the  car- 
bonates of  the  other  three  metals  of  this  group  ? 

Are  the  carbonates  of  calcium,  strontium  and  barium  soluble  in  acetic 
acid? 

State  how  barium  may  be  separated  from  calcium  and  strontium. 
How  may  the  presence  of  calcium  and  strontium  in  a  solution  of  their 
compounds  be  proven? 

(c)  If  a  solution  of  the  compounds  of  the  four  metals  of  this  group  were 
given  you,  how  would  you  proceed  to  prove  their  presence?      Make  a 
statement  of  your  proposed  method  and  submit  it  to  the  instructor  for 
inspection. 

(d)  Procure  an  "unknown"  solution  containing  two  or  more  of  the 
elements  of  this  group  and  make  an  analysis  of   same  according  to  the 
method  proposed  in  (c). 

RADIUM,    Rd. 

At.  Wt.  225          Sp.   Gr.   (?) 
(See  lecture  notes  and  references.) 

EXERCISES. 

i.  Chemically,  what  is  (a)  asbestos?  (b)  meerschaum?  (c)  olivine? 
(d)  serpentine? 


ALKALINE  EARTH  METALS.  315 

2.  What  is   (a)   chalk?  (b)   calcite?     (c)   limestone?     (d)   cement 
(Portland)  ?     (e)  apatite  ? 

3.  What  is  (a)  celestite?  (b)  strontianite  ? 

4.  What  is  (a)  witherite  ?  (b)  heavy  spar  or  barite  ? 

PROBLEMS. 

1.  Calculate  the  percentage  of  barium  in  barium  chloride. 

2.  Calculate  the  weight  of  sulphuric  acid  necessary  to  precipitate  the 
barium  as  barium  sulphate  from  a  solution  containing  .5  gram  of  barium 
chloride,  BaCl2.2H2O? 

3.  Calculate  the  percentage  of  chlorine  in  barium  chloride. 

4.  Calculate  the  weight  of  silver  nitrate  required  to  precipitate  the 
chlorine  as  silver  chloride  from  a  solution  containing  .7  gram  of  barium 
chloride. 


CHAPTER  XXVII. 

Zinc,  Zn,         65.4 

Cadmium,     Cd,       112.4 
Mercury,       Hg,       200.0 

Zinc,  cadmium  and  mercury  form  the  secondary  family  of  the  alkaline 
earths.  The  elementary  substances  of  this  family  and  their  corre- 
sponding compounds  possess  many  similar  properties.  Their  relation 
to  magnesium,  calcium,  strontium  and  barium  resembles  the  relation- 
ship existing  between  the  alkalies  and  copper,  silver,  and  gold. 

These  three  elementary  substances  are  lustrous  metals  of  high  specific 
gravity  and  are  found  in  combination  as  minerals,  which  occur  usually 
as  ores  in  the  older  crystalline  rocks.  Smithsonite,  ZnCO3,  sphalerite 
or  zinc  blende,  ZnS,  and  calamine,  H2Zn2SiO5,  are  the  principal  minerals 
which  contain  zinc.  Cadmium  is  found  in  small  quantities  in  those 
ores  in  which  zinc  is  the  chief  metal  present.  It  is  also  found  in  the 
rare  mineral,  greenockite,  CdS.  Small  quantities  of  mercury  are  found 
in  minute  globules  disseminated  through  the  pores  of  those  rocks  where 
the  most  important  mineral  containing  mercury,  namely,  the  red  sulphide, 
cinnabar  (HgS),  occurs.  Alloys  of  silver  and  gold,  known  as  amalgams, 
are  also  found  native.  The  metallurgy  of  these  substances  is  comparatively 
simple — all  three  of  the  metals  being  obtained  easily  by  roasting  the  ores 
alone  or  in  mixture  with  carbon,  in  ovens,  and  condensing  the  vapors 
as  the  metals  volatilize  at  high  temperatures. 

In  their  physical  characteristics,  the  elements  of  this  family  are  cer- 
tainly metallic  in  their  nature.  Zinc,  which  has  the  smallest  atomic 
weight  of  the  three,  is  less  positive  than  the  other  two.  The  appended 
table  shows  the  interesting  gradation  of  properties,  physical  and  chemical, 
with  increasing  atomic  weight;  for  example,  the  boiling  points  and  the 
melting  points  decrease  with  increase  of  atomic  weight.  It  will  be 
remembered  that  the  non-metallic  elements  behave  in  a  contrary  manner 
and  that  the  conduct  of  the  alkali  metals  is  very  similar. 

Zinc  is  a  bluish-white  crystalline  metal,  brittle  at  ordinary  tempera- 
tures, but  malleable  at  i2o°-i5o°.  Cadmium  is  a  white,  lustrous  metal, 
fairly  soft,  ductile  and  malleable.  Mercury  is  a  heavy,  silvery,  lustrous 
liquid. 

At  ordinary  temperatures  zinc  is  acted  upon  but  superficially  by  moist 
air,  but  cadmium  is  scarcely  affected  and  mercury  not  at  all.  However, 
all  three  of  the  metals  are  converted  into  oxides  when  heated  in  the  air. 

Finely  divided  zinc  will  set  free  hydrogen  from  water  at  ordinary 
temperature,  but  massive  zinc  and  cadmium  act  upon  water  only  at  red 
heat.  Mercury  has  no  effect  upon  water. 


ZINC.  317 

The  metals  of  this  family  can  be  deposited  from  solutions  of  their 
salts  by  the  electric  current  or  by  metals  above  them  in  the  electro- 
chemical series,  and  they  will  displace  other  metals  .below  them  in  the 
series. 

These  metals  are  bivalent — mercury  being  the  only  one  which  forms 
two  series  of  salts  (mercurous  and  mercuric  compounds),  which  differ 
from  each  other  in  the  proportions  in  which  mercury  is  contained  in 
them.  The  same  tendency  is  shown  to  a  slight  extent  by  cadmium. 

Zinc  dissolves  easily  in  dilute  acids,  and  is  attacked  by  the  caustic 
alkalies  when  the  latter  are  warmed.  Hydrogen  is  evolved  in  both  cases. 
Cadmium  dissolves  less  readily  in  acids,  but  is  not  soluble  in  solutions 
of  the  caustic  alkalies.  Mercury  does  not  dissolve  in  hydrochloric  acid 
or  dilute  sulphuric  acid;  concentrated  sulphuric  acid,  when  hot,  however, 
attacks  it  and  liberates  sulphur  dioxide;  concentrated  nitric  acid  acts 
upon  the  metal  forming  mercuric  nitrate;  dilute  nitric  acid  yields  mer- 
curous  nitrate.  The  caustic  alkalies  do  not  attack  mercury. 

The  oxides  and  hydroxides  of  zinc,  cadmium  and  mercury  are  soluble 
in  acids,  yielding  salts,  in  solutions  of  which  the  metals  are  positive 
ions.  A  few  zinc  salts,  known  as  zincates  and  of  the  type  Na2ZnO2,  in 
which  zinc  is  in  the  negative  ion,  are  formed  by  the  interaction  of  zinc, 
its  oxide  or  hydroxide,  with  strong  basic  hydroxides. 

2NaOH  +  Zn  — >  Na-jZnO-s  +~H2 
2NaOH  +  Zn(OH)2  —  Na2ZnO2  +  2H2O. 

In  this  respect  zinc  resembles  aluminum,  lead  and  a  number  of  other 
metals  which  can  display  both  metallic  and  non-metallic  properties. 

The  chlorides,  nitrates  and  sulphates  are  soluble  in  water,  but  the 
oxides,  carbonates  and  phosphates  are  insoluble  in  water.  The  chlorides 
are  comparatively  volatile  and  the  hydroxides  lose  water  forming  the 
oxides. 

The  compounds  of  zinc,  cadmium  and  mercury  do  not  give  color  to 
the  borax  bead. 

The  following  table  emphasizes  the  similarity  in  the  properties  of  the 
three  metals  and  their  corresponding  derivatives. 

Zinc.  Cadmium.           Mercury. 

Atomic  Weight,  65.4  112.4  200.0 

Specific  Gravity,  7.1  8.6                    13-59,  (o°) 

Atomic  Volume,  9.1  12.9                    14.7 

Melting-point,  419°  321.7°  —38.85° 

Boiling-point,  920°  (?)  778°  357° 

*Vapor  Density,  65.5  112  200.3 

*  These  vapor  density  determinations  indicate  that  the  molecular  weight  and  the 
atomic  weight  of  each  of  the.  three  elementary  substances,  when  they  are  in  the  state 
of  vapor,  are  identical.  If  our  conclusions,  as  to  the  relative  weights  of  these  atoms, 
are  correct,  it  is  obvious  that  zinc,  cadmium  and  mercury,  on  changing  from  the 
liquid  to  the  gaseous  state,  separate  into  individual  atoms.  Sodium  and  potassium 
behave  in  a  similar  manner. 


ZnO,H20 
82,600  cal. 

CdO,H2O 
65,700  cal. 

HgO 

22,000  cal 

ZnCl2 
97,200  cal. 

CdCl2 
93,200  cal. 

HgCl2 

54,500  cal. 

ZnS 
(Zn,S,Aq), 
39,600  cal. 

CdS 
(Cd,  S,  Aq.), 
32,400  cal. 

HgS 
Hg,S, 
16,200  cal. 

ZnSO4 
230.000  cal. 

CdSO4 
221.000  cal. 

HgS04. 

318  EXPERIMENTAL  CHEMISTRY. 

Oxides, 
H.  of  F. 

Chlorides, 
H.  of  F. 

Sulphides, 
H.  of  F. 

Sulphates, 
H.  of  F. 

Note. — It  is  maintained  by  many  that  if  we  were  able  to  make  vapor 
density  determinations  at  very  high  temperatures,  we  should  discover 
that  all  diatomic  molecules  are  converted  into  monatomic  ones. 

ZINC,   zn. 

At.  Wt.  65.4         Sp.  Gr.  7.1. 
Experiment  I. — Properties  of  Zinc. 

(a)  Clean  part  of  a  piece  of  zinc  with  sand-paper  or  a  file.     Color  ? 

(b)  Is  zinc  hard  or  soft  (use  the  point  of  a  knife  blade)  ? 

(c)  Does  zinc  tarnish  when  exposed  to  the  moist  air  ? 

(d)  Is  zinc  soluble  in  sulphuric  acid  ?     In  hydrochloric  acid  ?     Equa- 
tions? 

(e)  Place  about  i  cm.3  of  zinc  dust  in  a  test  tube,  and  add  5  cm.3  of  a 
strong  solution  of  sodium  hydroxide.     Heat  the  mixture  carefully  and 
test  the  evolved  gas  with  a  flame.     Results  ?     Identify  the  gas.     The 
solution    contains  sodium  zincate,  Na2ZnO2.     Express  the  action  by  an 
equation. 

(/)  Recall  or  repeat  Exp.  XI  (a),  "  Electrolysis  and  Electrical  Equiva- 
lents." 

(g)  Enumerate  some  of  the  uses  of  zinc.  What  is  "galvanized" 
iron? 

Experiment  II. — (Quant.)  Determination  of  the  Atomic  Weight  of 
Zinc. 

(a)  Proceed  as  in  Exp.  XIII,  "Hydrogen."     Calculate  the  weight  of 
hydrogen  collected.     Determine  the  atomic  weight  of  zinc  referred  to 
H  =  i,  remembering  that  zinc  is  bivalent.     Write  the  reaction. 

(b)  The  specific  heat  of  zinc  is  0.093.     Calculate  the  atomic  weight 
of  zinc  by  Dulong  and  Petit's  Law.     State  the  latter. 

Experiment  III. — Preparation  of  Hydrated  Zinc  Chloride. 

Dissolve  two  or  three  small  pieces  of  "granulated"  zinc  in  hydro- 
chloric acid,  dilute  with  an  equal  volume  of  water,  filter  and  evaporate 


ZINC.  319 

the  filtrate  to  oily  consistency  in  a  porcelain  dish.     Cool,  and  examine 
the  product  (?).     Equation? 

Experiment  IV. — Preparation  of  Hydrated  and  Dehydrated  Zinc  Sul- 
phate. 

(a)  Suggest  a  method  for  the  preparation  of  hydrated  zinc  sulphate, 
ZnSO4 .  yH2O.     Equation  ? 

(b)  How  may  the  above  salt  be  dehydrated?     At  what  temperature 
does  it  lose  all  or  a  part  of  its  water  of  hydration.    (Consult  your  text- 
book) ?     What  is  white  vitnol? 

Experiment  V. — Precipitation  and  Behavior  of  Zinc  Hydroxide.    Form- 
ation of  a  Zincate. 

(a)  To  a  dilute  solution  of  zinc  sulphate  add  a  sodium  hydroxide 
solution  drop  by  drop  and  shake.      Result?      What,  probably,  is  the 
precipitate?     Equation?     Filter,  and  suspend  the  precipitate  in   water 
in  a  test  tube  by  punching  a  hole  through  the  apex  of  the  filter  paper, 
and  directing  a  stream  of   water  from  the  water-bottle  upon   the  pre- 
cipitate.    Divide  the  mixture  into  two  parts. 

(b)  To  one  part  add  hydrochloric  acid?     Equation?     Does  the  zinc 
hydroxide  behave  in  this  reaction  as  a  base  or  an  acid  ? 

(c)  To  the  second  part  of  the  mixture  from  (a)  add  an  excess  of  the 
sodium  hydroxide  solution  (?).     The  compound  formed  by  the  inter- 
action   of   zinc    hydroxide    and  sodium    hydroxide    is    sodium    zincate, 
Na2ZnO2.     Ionic   equation?     Does  this   reaction  indicate  that  zinc  hy- 
droxide possesses  acidic  or  basic  properties?     Is  the  zinc  a  part  of  the 
anion  or  cation  ? 

What  are  your  inferences  as  to  the  "strength"  of  zinc  hydroxide  as  a 
base  ? 

Experiment  VI. — Preparation  and  Properties  of  Zinc  Hydroxycarbon- 
ate  and  Zinc  Oxide. 

(a)  Add  slowly  a  solution  of  sodium  carbonate  to  a  dilute  solution  of 
zinc  sulphate  ( ?).     Note  the  evolution  of  gas  as  the  precipitate  forms. 
Prove  that  the  escaping  gas  is  carbon  dioxide  (?).     Boil  the  mixture, 
taking  care  that  the  precipitant  is  in  excess;  filter,  and  wash  the  pre- 
cipitate.    Test  a  small  portion  of  the  precipitate  with  hydrochloric  acid. 
Is  carbon  dioxide  evolved  ?     Account  for  the  evolution  of  this  gas  during 
the  formation  of  the  precipitate.     Express  by  ionic  equations  the  pre- 
cipitation of  the  basic  zinc  carbonate. 

(b)  Heat  the  other  portion  of  the  precipitate  from  (a)  to  redness  in  a 
porcelain  crucible  until  a  portion  removed  and  tested  with  hydrochloric 
acid  does  not  effervesce.      Note  the  color  of  the  residue  when  hot  and 
when  cold  (?).     Equation?     Reserve  the  zinc  oxide  for  Exp.  VIII  (a). 


320  EXPERIMENTAL  CHEMISTRY. 

Experiment  VII. — An  Experimental  Study  of  Ionic  Equilibrium  and 
"Concentration  Effect." 

(a)  Place  5  cm.3  of  zinc  acetate,  zinc  sulphate  and  zinc  chloride  in 
separate  test  tubes.     Test  each  solution  with  litmus  paper  ( ?).     Saturate 
the    solutions    with    hydrogen    sulphide.     Results?     Equations?     Are 
these  reactions  reversible  theoretically;  i.e.,  does  the  acid  which  is  formed 
simultaneously  with  the  zinc  sulphide  in  each  case  dissolve  the  latter  to  a 
larger  or  smaller  degree?     Will  the  amount  of  zinc  sulphide  dissolved 
vary  with  the  activity  of  the  different  acids?     The  answers  to  these 
latter  questions  may  be  ascertained  by  two  different  experimental  methods. 
Proceed  as  follows: 

Filter  separately  the  contents  of  the  tubes  and  preserve  each  filtrate 
in  a  separate  tube  which  should  be  marked  to  prevent  confusion. 

Compare  the  rates  at  which  zinc  sulphide  (the  precipitate)  dissolves 
in  dilute  acetic  acid,  dilute  sulphuric  acid  and  dilute  hydrochloric  acid. 
Results?  What  are  your  inferences  as  to  the  relative  "activities"  of  the 
three  acids? 

Confirm  your  inferences  by  adding  ammonium  hydroxide  to  each 
filtrate  in  the  marked  tubes.  The  zinc  in  solution  is  now  precipitated 
as  zinc  sulphide  (the  ammonium  hydroxide  merely  neutralizes  the  free 
acid).  In  which  solution  (filtrate)  do  you  find  the  largest  amount  of 
zinc  ?  The  least  amount  ?  Do  these  results  confirm  your  previous 
conclusions  as  to  the  relative  strength  (activity)  of  the  three  acids  ? 

(b)  To  5  cm. 3  of  a  dilute  solution  of  zinc  sulphate  add  a  few  drops 
of   sulphuric    acid.     Pass   hydrogen    sulphide   into   the    solution.     If   a 
precipitate  forms,  add  more  acid,  drop  by  drop,  until  hydrogen  sulphide 
does    not    produce    a    precipitate.       The    addition   of   sulphuric    acid 
increases   the   concentration   of   what   ions?     Does   this   influence    the 
degree  of  ionization  of  the  hydrogen  sulphide  ?     When  there  is  an  excess 
of  free  acid  present  in  the  solution  is  the  hydrogen  sulphide  chiefly  in 
the  molecular  condition  or  is  it  largely  dissociated  ?      Is  chemical  action 
an  interaction  of  ions  or  molecules  ordinarily  ?      Show  by  ionic  equation 
the  failure  of  hydrogen  sulphide  to  precipitate  zinc  sulphide  in  the  presence 
of  a  free  inorganic  acid.     Now  add  an  excess  of  a  sodium  acetate  solution 
(zinc  acetate  is  soluble),  and  pass  hydrogen  sulphide  into  the  solution 
if  a  precipitate  of  zinc  sulphide  does  not  form  at  once.     Account  for  the 
formation  of  the  precipitate.     Express  the  various  interactions  by  ionic 
equations. 

Experiment  VIII. — Analytical  Reactions  of  Zinc  Salts. 

(a)  Heat  a  small  piece  of  zinc  on  charcoal  in  the  oxidizing  flame  ( ?). 
Moisten  the  incrustation  which  is  formed  on  the  charcoal  with  a  drop  of 
cobalt  nitrate,  and  heat  again  or  place  a  small  pinch  of  the  zinc  oxide 
prepared  in  Exp.  VI  (b)  on  the  charcoal  and  proceed  as  above.  Result  ? 
Equations?  (The  formula  of  the  green  compound,  Rinmaris  green,  is 
CoZnO2.) 


CADMIUM.  321 

Does  the  zinc  manifest  acidic  or  basic  properties  in  this  reaction? 
Reasons  for  your  answer? 

(b)  Test  a  dilute  solution  of  zinc  sulphate  with  litmus  paper  (?). 
Account  for  the  result. 

(c)  Ammonium  hydroxide  produces  in  neutral  solutions  of  zinc  salts 
a   partial   precipitation   of    zinc   hydroxide   which   is   soluble   in  excess 
of  the  precipitant    and  ammonium  chloride,  forming  Zn(NH3)4.Cl2  or 
Z  n(NH3).4(OH)2.     If  this  solution  is  treated  with  ammonium  sulphide, 
(NH4)2S,  zinc  sulphide  is  formed.     Equations? 

(d)  Add  a  slight  excess  of  ammonium  sulphide  to  several  cm.3  of  a 
solution  of  zinc  sulphate.     Result  ?     Equation  ?     Filter,  and  wash  the 
precipitate   with   water.     Divide   the   precipitate   into   two   parts.     Try 
its  solubility  in  dilute  hydrochloric  acid,  and  acetic  acid  ( ?).     Equations? 

Why  was  ammonium  sulphide  used  instead  of  hydrogen  sulphide 
in  the  above  reaction  ? 

What  is  the  color  of  the  zinc  ion  ?     Its  valence  ? 

CADMIUM,    Cd. 

At.  Wt.   112.4         Sp.   Gr.  8.6. 

Experiment  I. — Properties  of  Cadmium. 

(a)  Examine  a  piece  of  the  metal  and  tabulate  its  obvious  physical 
and  chemical  properties. 

(b)  Will  cadmium  displace  zinc  from  a  solution  of  a  salt  of  the  latter 
(Note  their  relative  positions  in  the  electro-chemical  series)  ? 

(c)  Calculate  the  atomic  weight  of  cadmium  by  Dulong  and  Petit's 
Law,  the  specific  heat  of  cadmium  being  0.054. 

(d)  Give  the  composition  of  "Wood's  Metal."  (See  Appendix).     Re- 
peat or  review  Exp.  II.  "Bismuth." 

(e)  Enumerate  some  of  the  chief  uses  of  the  metal. 

Experiment  II. — Properties  of  Cadmium  Salts. 

Same  as  Exp.  II  "Magnesium." 

Experiment  III. — Precipitation  and  Solubility  of  Cadmium  Hydroxide. 

Repeat  Exp.  V  (a)  and  (b),  substituting  the  word  cadmium  where 
the  word  zinc  appears.  Equations?  Compare  the  behavior  of  the  zinc 
and  the  cadimum  compounds  (?). 

Experiment  IV. — Precipitation  of  Cadmium  Carbonate.  Preparation 
of  Cadmium  Oxide. 

(a)  To  10  cm.3  of  a  dilute  solution  of  cadmium  chloride  add  a  small 
excess   of   sodium   carbonate.     Result  ?     Equation  ?     Filter,    and    wash 
the  precipitate.     Test  a  small  portion  of  the  precipitate  with  hydrochloric 
acid  (?).     Equation? 

(b)  Heat  the  rest  of  the  precipitate  from  (a)  to  redness  in  a  crucible 


322  EXPERIMENTAL  CHEMISTRY. 

until  a  portion  removed  and  tested  with  hydrochloric  acid  does  not 
effervesce.  Note  the  color  of  the  residue  when  hot  and  when  cold  (?). 
Equations?  Reserve  the  cadmium  oxide  for  Exp.  V  (a). 

Experiment  V. — Analytical  Reactions  of  Cadmium  Salts. 

(a)  Mix  a  portion  of  the  cadmium  oxide  obtained  from  Exp.  IV  (b),  or 
some   other  cadmium   compound,   with   anhydrous   sodium   carbonate, 
and  heat  the  mixture  on  a  piece  of  charcoal  in  the  reducing  flame.     Dur- 
ing the  preliminary  stages  of  the  reduction,  note  the  characteristic  in- 
crustation (?).     Continue  the  process  until  small  metallic  globules  (?) 
make  their  appearance. 

(b)  Same  as  Exp.  VIII  (b)  "Zinc." 

(c)  Ammonium    hydroxide    precipitates    cadmium    hydroxide    from 
solutions  of  cadmium  salts.     The  hydroxide  is  soluble  in  excess,  form- 
ing salts  such  as  Cd(NH3)4.SO4  or  Cd(NH3)4.  (OH)2.     If  this  ammo- 
niacal  solution  is  treated  with  a  solution  of  potassium  cyanide,  KCN,  a 
soluble  complex  salt,  K2.Cd(CN)4  is  formed. 

(d)  Saturate  a  solution  of  cadmium  chloride  with  hydrogen  sulphide  ( ?) . 
Equation?      Is    the    reaction    reversible?       Answer    this    question 
by  testing  the  solubility  of  cadmium  sulphide  in  dilute  hydrochloric  acid. 
Will  hydrogen  sulphide  precipitate  cadmium  sulphide  from  solutions 
of  cadmium  salts  containing  a  slight  amount  of  free  acid?     Recall  the 
behavior  of  zinc  salts  in  the  presence  of  the  same  reagents  ( ?). 

Test  the  solubility  of  cadmium  sulphide  in  concentrated  nitric  acid  ( ?). 
Equation?  In  ammonium  sulphide  (?). 

(e)  To  a  solution  of  cadmium  chloride  add  a  solution  of  potassium 
cyanide  (Care!  Poison!),  drop  by  drop.     Observe  the  precipitation  of 
cadmium  cyanide,   Cd(CN)2.     Equation?     Add  an  excess  of  the  pre- 
cipitant (?).     The  soluble  complex  salt  formed  is  K2.Cd(CN)4.    Equa- 
tion?    Is  the  cadmium  ion  present  in  the  solution,  theoretically?  (See 
discussion    of    "double"    and     " complex"    salts  under   Exp.   XVII, 
"Copper".)     Confirm  or  disprove  your  conclusions  by  passing  hydrogen 
sulphide    into    the    above    solution.     Results?     Ionic    equations?     Is 
cadmium  sulphide  soluble  in  potassium  cyanide  ? 

(/)  What  is  the  color  and  valence  of  the  cadmium  ion  ? 

(g)  State  the  tests  by  which  you  could  distinguish  between  the  salts 
of  zinc  and  cadmium.  Apply  to  the  assistant  for  an  "  unknown  "  solution. 
Ascertain  if  the  solution  contains  one  or  both  of  the  above  metals.  Make 
a  complete  report  of  your  procedure. 

MERCURY,    Hg. 

At.  Wt.   200.0         Sp.   Gr.   13.59. 
Experiment  I. — Properties  of  Mercury. 

(a)  Tabulate  the  most  obvious  physical  properties  of  mercury.  What 
particular  property  does  it  possess  which  makes  it  unique  among  metals  ? 


MERCURY.  323 

(b)  Does  mercury  tarnish  when  exposed  to  the  air? 

(c)  Enumerate  those  properties  of  mercury  in  virtue  of  which  it  appeals 
to  you  as  a  metal. 

(d)  Is  the  molecule  of  mercury  monatomic  or  diatomic  ? 

(e)  State  some  of  the  uses  of  mercury.     What  are  amalgams  ? 

Experiment  II. — Preparation  of  Mercury  by  Roasting  Cinnabar. 

Place  a  small  piece  of  cinnabar  (HgS)  in  a  glass  tube  open  at  both 
ends;  clamp  the  tube  in  an  inclined  position  so  as  to  permit  a  free  draught 
of  air  through  it.  Heat  the  tube  strongly  with  the  Bunsen  flame.  Results? 
Identify  the  fumes  ( ?).  Equation  ? 

Experiment  III. — Properties  of  the  Salts  of  Mercury. 

Examine  the  different  salts  of  mercury  (end  shelf),  noting  the  obvious 
physical  properties.  Heat  a  small  quantity  of  a  mercury  salt  in  a  test 
tube.  Result?  This  property  of  volatilizing  unchanged,  giving  sub- 
limates of  the  same  compound,  is  characteristic  of  many  mercury  com- 
pounds. Give  the  name,  formula  and  color  of  each  salt.  How  many 
series  of  mercury  salts  do  you  find  ?  Tabulate  the  foregoing  data. 

Give  the  formula  and  use  of  each  of  the  following  substances:  calo- 
mel, corrosive  sublimate,  vermillion. 

Experiment  IV. — Formation  of  Mercurous  Nitrate  from  Dilute  Nitric 
Acid  and  Mercury. 

By  means  of  a  glass  pipette  take  two  or  three  drops  of  mercury  from 
the  supply-bottle,  place  the  mercury  in  a  test  tube  or  beaker,  and  add 
15  cm.3  of  dilute  nitric  acid  (i  to  i).  Allow  the  action  to  continue  for 
about  an  hour  (heating  gently  will  hasten  the  action).  An  excess  of 
mercury  must  always  be  present.  Continued  stirring  will  cause  crystal- 
lization of  the  mercurous  nitrate  which  is  formed  in  the  solution  under 
the  above  conditions.  Dissolve  the  crystals  in  distilled  water  which  has 
been  made  slightly  acid  by  the  addition  of  a  few  drops  of  nitric  acid. 
Filter  the  solution,  and  reserve  the  filtrate  for  use  in  those  experiments 
in  which  mercurous  nitrate  is  demanded.  Equation?  What  is  the 
valence  of  mercury  in  mercurous  nitrate  ? 

Experiment  V. — Formation  of  Mercuric  Nitrate  from  Concentrated 
Nitric  Acid  and  Mercury. 

Boil  10  cm.3  of  mercurous  nitrate  or  two  or  three  globules  of  mercury 
with  an  excess  of  strong  nitric  acid.  Continue  the  boiling  until  all 
of  the  mercury  has  disappeared,  evaporate  nearly  to  dryness  upon  a 
steam  bath,  add  a  few  drops  of  nitric  acid,  and  continue  the  evaporating 
until  the  reddish-brown  fumes  cease  to  be  liberated.  The  residue  is 
mercuric  nitrate.  Prepare  a  solution  of  the  salt  according  to  the  direc- 
tions given  in  the  preceding  experiment.  Reserve  the  solution  for 
experiments. 


324  EXPERIMENTAL  CHEMISTRY. 

Experiment  VI. — Precipitation  of  Mercurous  Iodide  and  Mercuric 
Iodide. 

(a)  Add  a  solution  of    potassium    iodide  drop  by  drop    to  a  dilute 
solution  of  mercurous  nitrate.     Result  ?     Equation  ?     What  is  the  effect 
when  an  excess  of  the  precipitant  is  added  ? 

(b)  Repeat  (a)  using  mercuric  nitrate  instead  of  mercurous  nitrate. 
Observe  the  precipitation  of  the  yellow  variety  of  mercuric  iodide,  which 
in  a  few  minutes  change  into  the  red  variety.     The  conversion  into  the 
red  form  is  greatly  hastened  by  light.     Add  an  excess  of  the  solution  of 
potassium  iodide.     Result  ?     What  complex  ion  is  formed  ?     Equation  ? 

Note. — When  the  red  iodide  of  mercury  is  heated  above  126°  it  is 
converted  into  the  yellow  crystalline  variety  of  mercuric  iodide  which  if 
kept  in  the  cold,  changes  slowly  into  the  red.  If  the  yellow  variety  is 
scratched  or  rubbed,  it  is  converted  at  once  into  the  red  crystalline  form. 
The  red  variety  is  stable  at  temperatures  below  126°,  and  the  yellow, 
above  126°.  This  is  the  transition  temperature  which  separates  the  two 
regions  of  stability.  Substances  like  sulphur  and  mercuric  iodide  which 
can  change  in  two  directions-i.e.,  have  two  regions  of  stability-are  said 
to  possess  the  property  of  "  enantiotropy."* 

Experiment  VII. — Analytical  Reactions  of  Mercury  Salts. 

(a)  Heat  a  small  quantity  of  a  salt  of  mercury  with  an  equal  amount 
of  anhydrous  sodium  carbonate  in  a  dry  test  tube.     Result  ? 

(b)  Recall     or    repeat    Exp.    XI  (d)  "  Electrolysis    and     Electrical 
Equivalents." 

Note. — To  separate  portions  of  a  dilute  solution  of  mercurous  nitrate 
add  the  following  reagents.  Repeat  each  experiment  using  a  dilute 
solution  of  any  mercuric  salt.  Compare  and  tabulate  the  results. 

(c)  Test  with  litmus  paper  ( ?).     Explain. 

(d)  A  solution  of  sodium  hydroxide  (?).     Equations? 

(e)  Ammonium  hydroxide  (?).     Equations? 

(/)  Dilute  hydrochloric  acid  (?).  If  a  precipitate  is  formed,  filter 
and  treat  the  precipitate  upon  the  filter  paper  with  ammonium  hydroxide. 
Equations  ? 

Note. — The  compound  formed  by  the  action  of  the  ammonium  hy- 
droxide is  mercurous  chloramide,  NH2.Hg2.Cl.  A  compound  formed 
by  the  action  of  ammonium  hydroxide  on  mercuric  chloride  is  known 
as  mercuric  chloramide,  NH2.Hg.Cl. 

(g)  Hydrogen  sulphide  ( ?).  If  a  precipitate  is  formed,  ascertain  its 
solubility  in  strong  nitric  acid.  Equation? 

(h)  Stannous  chloride  (?).  Use  an  excess  of  the  reagent  and  warm 
gently  (?).  Equations? 

*  Ostwald — Principles  of  Inorganic  Chemistry. 


MERCURY.  325 

(i)  Compare  the  properties  of  the  monomercurion,  Hg',  and  the  dimer- 
curion,  Hg". 

(/)  State  how  a  solution  of  a  mercurous  compound  can  be  distin- 
guished from  a  solution  of  a  mercuric  salt. 

How  could  you  separate  the  metals  from  a  solution  containing  the 
salts  of  zinc,  cadmium  and  mercury?  Make  a  brief  coherent  report 
of  the  method  of  procedure. 

PROBLEMS. 

1.  The  atomic  weight  of  mercury  is  200.0.      If  the  vapor  density  of 
mercury  is  100,  how  many  atoms  does  the  molecule  contain  ? 

2.  10  grams  of  zinc  will  liberate  from  sulphuric  acid  how  many  cm.3 
of  hydrogen  at  o°  C.  and  760  mm.  ?     At  20°  C.,  and  730  mm.  ? 

3.  Calculate  the  percentage  composition  of  cadmium  sulphide. 


CHAPTER  XXVIII. 
THE  ELEMENTS  OF  GROUP  III. 

Family  M.  Family  m. 

(Scandium,       Sc.      44.  )  Boron,         B.       n.o 

(Yttrium,          Y.       89.0)  Aluminum,  Al.     27.1 

(Lanthanum,    La.   138.9)  (Gallium,    Ga.     70.0) 

(Ytterbium,       Yb.  173.0)  (Indium,      In.    115.  ) 

Thallium,    Tl.    204.1 

Scandium,  yttrium,  lanthanum  and  ytterbium,  like  certain  other 
elements  of  the  following  groups,  are  generally  known  as  rare  earth 
metals.  Cerium  (Ce.;  At.  Wt.i4o.2),  praseodymium  (Pr.;  At.  VVt.  140.5), 
neodymium  (Nd.;  At.  Wt.  143.6),  samarium  (Sa.;  At.  Wt.  150.3),  gado- 
linium (Gd.;  At.  Wt.  156),  and  erbium  (Er.;  At.  Wt.  166)  are  usually  in- 
cluded amongst  the  metals  of  the  rare  earths.  They  have  been  given 
this  generic  name  because  they  are  found  only  in  small  amounts  in  a  few 
rare  earthy  minerals,  such  as  orthite,  euxenite,  cerite,  gadolinite  and 
monazite.  Very  few  of  these  metals  have  been  isolated  in  pure  elementary 
form,  but  many  of  their  salts  have  been  prepared.  The  latter,  however, 
are  so  similar  in  behavior  that  separation  is  exceedingly  difficult. 

The  study  and  isolation  of  the  elements  of  this  group  is  attended 
by  much  difficulty.  This  is  well  illustrated  by  the  history  of  the  isolation 
of  the  compounds  of  the  elements  of  praseodymium  and  neodymium. 
In  1839,  Mosander  prepared  from  the  mineral  yttria,  certain  oxides, 
among  which  there  was  one  which  he  regarded  as  a  compound  of  a  new 
elementary  substance,  to  which  the  name  didymium  was  eventually 
given.  Nearly  sixty  years  elapsed  before  the  elementary  character  of 
didymium  was  questioned.  The  comparatively  recent  researches  of 
Auer  von  Welsbach  have  shown,  however,  that  didymium  is  a  mixture 
of  the  two  elements,  praseodymium  and  neodymium.  The  salts  of 
praseodymium  are  light  green  while  the  neodymium  salts  possess  a 
rose-violet  color. 

The  properties  of  these  elements  are  similar  to  those  of  aluminum, 
showing  the  same  gradations  as  wrere  found  in  preceding  groups,  with 
increase  of  atomic  weights;  that  is,  the  free  metals  possessing  the  greater 
combining  weights  are  the  more  readily  oxidized  and  their  respective 
bases  are  stronger.  In  their  compounds  the  metals  are  usually  trivalent. 
The  oxides  and  hydroxides  are  insoluble  in  water,  and  the  chlorides, 
nitrates  and  sulphates  are  soluble.  The  nitrates  are  decomposed  by 
heat.  "Many  of  these  elements  give  a  very  complicated  spectrum  on 
allowing  the  electric  spark  to  pass  between  carbon  points  moistened  with 
solutions  of  their  salts.  Since,  under  given  conditions,  each  element 


THE    ELEMENTS    OF    GROUP    III. 


327 


possesses  a  perfectly  definite  spectrum,  it  can  be  seen  whether  the  spec- 
trum changes  by  partial  separations.  Where  this  is  the  case,  we  are 
certainly  dealing  with  a  mixture.  The  higher  members  also  exhibit 
absorption  spectra,  some  of  them  also  emission  spectra." 

Several  of  the  metals  of  this  group  have  atomic  weights  so  near  to 
lanthanum  that  it  seems  impossible  at  present  to  accommodate  them 
in  the  periodic  table.  Ostwald.*  in  discussing  the  probable  position  of 
lanthanum  in  the  table,  says:  "  This  signifies  that  there  exist  at  this  point 
not  one  element,  but  a  number  of  elements  which  are  all  very  close  to  one 
another,  and  have  therefore  an  almost  equal  claim  to  this  position.  This 
is  an  occurrence  of  numerous  small  planetary  bodies  at  a  point  of  the 
solar  system  where,  by  analogy,  one  would  have  expected  a  large  planet." 

Scandium,  which  was  discovered  in  1879  by  Nilson  and  Cleve,  is  of 
particular  interest  to  us  because  its  existence  was  suggested  in  1869  by 
Mendeleeff  who,  from  considerations  based  upon  the  periodic  table, 
predicted  its  atomic  weight  and  many  of  its  physical  and  chemical 
properties.  Mendeleeff  had  given  the  name  of  eka-boron  to  the  element. 


Eka-boron  (Predicted  1869). 

Atomic  weight,  about  44. 

Oxide,  Eb2O3,  soluble  in  acids, 
analogous  to  A12O3,  but  more 
basic;  insoluble  in  alkalies. 


Scandium  (Discovered  1879). 

Atomic  weight,  44.1. 

Oxide,  Sc2O3,  soluble  in  strong 
acids,  analogous  to  A12O3,  but 
much  more  basic;  insoluble  in 
alkalies. 


Salts    colorless    and    give    gelatin-     Salts  colorless  and  give  gelatinous 
ous    precipitates  with  NaOH  or         precipitates    with     NaOH     or 

Na2CO3. 


Sulphate,  Sc2(SO4)3,  forms  a 
double  salt  with  K2SO4,  which  is 
not  isomorphous  with  the  alums. 

Gallium  (Discovered  1875). 
Atomic  weight,  69.9. 
Melting-point,  30.2°. 
Specific  gravity,  5.93. 

Slightly  oxidized  at  red  heat. 
Decomposes   water   at   high   tem- 
perature. 

Gallium  oxide,  Ga2O3. 
Gallium  chloride,  Ga2Cl6. 
Gallium  sulphate,  Ga2(SO4)3. 
Forms  a  well-defined  alum. 


Sulphate,  Eb2(SO4)3,  will  form  a 
double  salt  with  K2SO4,  not  iso- 
morphous  with  the  alums. 

Eka-aluminum  (Predicted  1871). 
Atomic  weight,  about  69. 
Will  have   a   low  melting-point. 
Specific  gravity,  about  5.9. 

Will  not  be  acted  upon  by  the  air. 
Will  decompose  water  at  red  heat. 

Will  give  an  oxide,  E12O3. 
Will  give  a  chloride,  E12O6. 
Will  give  a  sulphate,  E12(SO4)3. 
Will  form  a  potassium  alum,  etc. 


*Principles  of  Inorganic  Chemistry. 


328  EXPERIMENTAL  CHEMISTRY. 

With  the  exception  of  boron,  aluminum  and  thallium,  the  members 
of  Family  m.  are  among  the  rarest  of  the  elements.  Gallium,  indium 
and  thallium  were  discovered  by  means  of  the  spectroscope;  gallium 
takes  its  name  from  the  country  (France)  in  which  it  was  discovered  by 
Lecoq.  de  Boisbaudrau,  in  1875;  indium,  by  Reich  and  Richter  in  1863, 
received  its  name  on  account  of  two  characteristic  lines  in  the  indigo- 
blue  part  of  the  spectrum;  and  thallium,  by  Crookes  in  1861,  owes  its 
name  to  the  fact  that  there  is  a  prominent  green  line  in  its  spectrum 
(Gk.  OaXXSf,  a  green  twig).  These  metals  occur  in  exceedingly  small 
quantities  as  impurities  in  zinc  blends.  They  are  moderately  heavy, 
similar  to  one  another  in  most  properties,  and  possess  specific  gravities 
and  melting  points  which  vary  in  the  order  of  their  atomic  weights. 

Gallium  is  a  lustrous  gray  metal  which  possesses  the  extremely  low 
fusing-point  of  30.2°  C.  It  is  oxidized  superficially  when  heated  to  high 
temperatures  in  the  air,  and  is  acted  upon  by  strong  acids  and  strong 
basic  hydroxides.  Indium  is  a  soft  white  metal  which  melts  at  176°. 
It  burns  with  a  violet  flame  when  heated  in  the  air,  and  is  acted  upon  by 
strong  acids.  Thallium  is  a  soft  grayish-white  metal  which  oxidizes 
at  ordinary  temperature  in  moist  air.  It  melts  at  290°.  Sulphuric  acid 
and  nitric  acid  act  upon  thallium,  but  hydrochloric  acid  acts  upon  it 
only  superficially  as  the  insoluble  thallous  chloride  which  forms  upon  the 
surface  protects  the  metal. 

The  hydroxides  of  these  three  elements  are  weak  bases.  Gallium 
hydroxide  and  indium  hydroxide  may  also  interact  with  strong  basic 
hydroxides,  but  thallium  hydroxide  shows  no  such  acid  properties. 
Gallium  functions  both  as  a  bivalent  and  a  trivalent  element;  indium, 
as  a  univalent,  bivalent  and  trivalent  element,  and  thallium,  as  a  univalent 
and  a  trivalent  element. 

Boron  is  the  only  non-metal  of  the  group,  all  the  others  exhibit  well 
marked  basic  properties.  The  oxide,  B2O3,  of  boron  is  acidic.  (See 
Chapter  XXII). 

Aluminum  does  not  occur  native,  but  its  oxides  and  silicates  are  found 
widely  distributed.  It  is  prepared  by  the  electrolysis  of  oxide  of  alumi- 
num dissolved  in  cryolite.  Aluminum  is  a  silver-white  metal,  very 
ductile  and  malleable,  a  good  conductor  of  electricity,  and  melts  at  657°. 
It  does  not  tarnish  and  is  practically  without  action  upon  water.  It  is 
scarcely  acted  upon  by  nitric  acid,  but  readily  dissolves  in  hydrochloric 
acid  and  in  strong  solutions  of  the  caustic  alkalies  with  liberation  of 
hydrogen. 

Aluminum  functions  as  a  trivalent  element  in  its  compounds.  It  is  a 
highly  electro-positive  element,  and  in  consequence  of  its  great  affinity 
for  oxygen,  aluminum  displaces  all  metals  save  magnesium  from  their 
oxides.  The  extreme  readiness  with  which  aluminum  is  able  to  effect 
such  reduction  when  the  reaction  is  once  started  by  heat  and  the  ex- 
ceedingly high  temperature  which  is  reached  by  the  action  have  led  to 
some  very  useful  applications.  The  heat  of  reaction  between  powdered 
aluminum  and  ferric  oxide  has  found  useful  application  in  the  "Gold- 


THE    ELEMENTS    OF    GROUP    III.  329 

schmidt  process"  for  welding  iron  and  steel.  The  very  high  temperature, 
25oo°-3ooo°,  produced  by  the  reaction  is  sufficient  to  melt  both  the 
iron  and  the  oxide  of  aluminum.  As  the  products  of  the  action  are  not 
miscible,  they  separate  into  two  layers.  The  sulphides  are  reduced 
just  as  readily.  This  furnishes  a  simple  method  for  the  preparation  of 
pure  specimens  of  the  metals  whose  oxides  and  sulphides  are  reduced 
with  difficulty.  Goldschmidt,  the  inventor,  has  given  to  these  processes 
the  general  name  of  "  ahiminothermy"  Mixtures  of  the  metallic  oxides 
and  granulated  aluminum  have  been  placed  on  the  market  under  the 
name  of  "thermit" 

The  hydroxide  of  aluminum  is  feebly  acidic  as  well  as  basic,  and 
therefore  forms  two  classes  of  compounds  of  the  types  Na3AlO3  and 
A1C13.  The  general  name  of  aluminates  is  given  to  the  first  class.  Both 
series  of  salts  are  hydrolyzed  by  water,  which  is  to  be  expected. 

Aluminum  has  a  wide  range  of  uses,  although  the  marked  influence 
of  traces  of  impurities  has  limited  its  application  more  than  was  anticipated 
originally.  It  is  used  for  making  aluminum  bronze,  magnalium,  flash- 
light powders,  paint,  cables  for  conducting  electricity,  foil,  various 
ornamental  articles  and  cooking  utensils. 

The  compounds  of  aluminum  are  of  equal  importance.  Ruby,  corun- 
dum and  sapphire  are  nearly  pure  oxides  of  aluminum.  Emery  is  a 
mixture  of  corundum  and  iron.  Clay,  an  impure  silicate  of  aluminum, 
is  used  in  the  manufacture  of  bricks  and  earthenware.  Kaolin,  nearly 
pure  aluminum  silicate,  is  used  in  the  manufacture  of  porcelain  and 
china.  Hydraulic  cement  (hardens  under  water)  is  made  by  heating  a 
mixture  of  limestone  and  clay  to  incipient  fusion. 

BORON,  B. 

At.  Wt.  n.o     Sp.  Gr.  2.5. 
(See  Chapter  XXII). 

ALUMINUM,  Al. 
At.  Wt.  27.1     Sp.  Gr.  2.58. 
Experiment  I. — Properties  of  Aluminum. 

(a)  Examine  specimens  of  aluminum  in  the  forms  of  ingot,  wire  and 
powder.     Introduce  a  piece  of  the  wire  into  the  Bunsen  flame  in  order 
to  determine  whether  aluminum  is  a  conductor  of  heat  (?).     Does  the 
wire  melt  in  the  flame  (use  tongs)?     Try  the  blast- lamp  (?).     What  is 
the  melting  point  of  the  metal?     Does  it  tarnish  in  the  air? 

(b)  Solubility  in  acids.     Try  the  action  of  the  powdered  or  granulated 
metal  on  the  following  reagents:  hydrochloric  acid  (?),  nitric  acid  ( ?), 
and  sulphuric  acid  (?).     Account  for  results.     Test  the  gas,  if  any  is 
liberated  (?).     Equations? 

(c)  Solubility  in  the  caustic  alkalies.     Add  about  5  cm. 3  of  a  con- 


330  EXPERIMENTAL  CHEMISTRY. 

centrated  solution  of  sodium  hydroxide  or  potassium  hydroxide  to  i  to  2 
cm.3  of  granulated  aluminum,  and  warm  gently  (?).  Identify  the  gas 
evolved  (?).  Add  more  of  the  hydroxide  if  necessary.  The  solution 
contains  sodium  aluminate,  Na3AlO3.  Ionic  equation  ?  Does  the  acidic 
or  basic  properties  of  the  aluminum  hydroxide  predominate  in  this 
reaction?  Your  reasons?  To  prove  that  aluminum  has  gone  into 
solution,  neutralize  the  product  of  the  above  reaction  carefully  with 
dilute  hydrochloric  acid.  Result  ?  Equation  ?  Are  the  properties  of  the 
aluminum  revealed  by  this  reaction  ?  Explain. 

(d)  Tabulate    the    properties    of    aluminum.     Should    commercial 
"lyes"  (alkalies)  be  warmed  in  cooking  utensils   made  of  aluminum? 
Why? 

(e)  Enumerate  some  of  the  uses  of  aluminum. 

Experiment  II.— Reduction  of  Metallic  Oxides  by  Use  of  Aluminum. 
The  Goldschmidt  Process.  "Thermit." 

Thoroughly  mix  small  quantities  of  iron  oxide  and  granulated  alumi- 
num. Place  the  mixture  on  an  iron  plate  or  a  piece  of  tile,  and  ignite  it 
by  means  of  a  piece  of  magnesium  ribbon  ( ?).  Examine  the  fused  mass 
for  globules  of  iron.  If  the  above  mixture  does  not  ignite  readily,  procure 
a  small  quantity  of  commercial  mixture  known  as  "thermit  iron" 
(iron  oxide  and  aluminum)  and  ignite  it  after  receiving  instructions  from 
the  assistant.  Account  for  the  results.  Equation  ? 

Experiment  III. — Salts  of  Aluminum. 

Examine  the  various  aluminum  salts  found  in  the  laboratory.  Record 
the  name,  formula  and  color  of  each  salt.  What  is  ultramarine,  its 
formula,  color  and  use? 

Experiment  IV. — Preparation  of  the  Compound  Salt,  Potassium- 
Aluminum  Sulphate.  Alums. 

(a)  Prepare  hot  saturated  solutions  of  potassium  sulphate  and  alumi- 
num  sulphate — 25  cm.3  of  each  solution  will  be  sufficient.     Mix  the 
two    solutions   in   a    small   beaker   or   crystallizing    dish   so    that   the 
resulting  solution  contains  the  weights  of  the  two  salts  in  a  proportion 
approximating    the   ratio   of    their    molecular   weights    (10    grams    of 
Al2(SO4)3.i8H2O  and  3  grams  of  K2SO4  in  200  cm.3  of  distilled  water. 
Concentrate    to   100    cm.3    Cool.)      The    compound    salt,    potassium- 
aluminum  sulphate,  K2SO4,   A12(SO4)3. 24H2O,  crystallizes  on  standing 
(it  may  be  necessary  to  allow  it  to  stand  for  several  days).     What  is  the 
form  of  the  crystals  ? 

(b)  Repeat  (a),  using  ammonium  sulphate  instead  of  potassium  sul- 
phate.    Result?     Compare  the  forms  of  the  crystals  (?). 

(c)  What  is  an  alum  ?     State  some  of  the  uses  of  alums.     Notice  the 
composition  of  the  various  alums,  a  list  of  which  is  appended.     All  of 
these  alums  crystallize  in  the  same  form  and  possess  the  same  funda- 
mental chemical  poperties. 


THE    ELEMENTS    OF    GROUP    III.  331 

Na2SO4.Al2(SO4)3.24H2O,        Sodium  alum. 
K2SO4.A12(SO4)3.24H2O,          Potassium  alum. 
(NH4)2SO4.  A12(SO4)3.24H2O,   Ammonium  alum. 

Alums  are  also  formed  by  substituting  iron,  chromium  and  man- 
ganese for  aluminum. 

K2SO4.Fe2  (SO4)3.24H2O,         Iron  alum. 
K2SO4 .  Cr2(SO4)3 . 24H2O,          "  Chrome  "  alum. 
K2SO4 .  Mn2(SO4)3 . 24H2O,         Manganese  alum. 

Experiment  V. — Analytical  Reactions  of  Aluminum  Salts. 

(a)  Heat  a  small  quantity  of  an  aluminum  salt  (use  alum)  on  a  piece 
of  charcoal  in  the  oxidizing  flame  (?),  moisten  with  a  few  drops  of  a 
cobalt  nitrate  solution,  and  heat  again.     Account  for    the  blue  color  of 
the  product. 

(b)  Test  a  solution  of  the  chloride  or  sulphate  of  aluminum  with  litmus 
paper  (?).     Explain. 

(c)  To  10  cm.3  of  an  aluminum  sulphate  solution  add  a  small  quantity 
of  a  solution  of  sodium  hydroxide.     Filter  off  the  precipitate,  and  suspend 
it 'in  water.     Divide  the  mixture  into  two  portions.     To  one  portion 
add  hydrochloric  acid  (?).     Equations?     To  the  second  portion  add  a 
slight  excess  of  the  sodium  hydroxide  solution  (?).     Now  add  an  excess 
of  a  solution  of  ammonium  chloride  ( ?).     Account  for  the  re-precipi- 
tation of  aluminum  hydroxide.     Equations? 

(d)  Add   a   sodium  carbonate   solution   to   a   solution   of   aluminum 
sulphate  (?).     Filter  off  the  precipitate,  and  wash  it  until  free  from 
sodium  carbonate.     Test  the  precipitate  with  hydrochloric  acid.     Is  it 
soluble  and  is  carbon  dioxide  evolved.     Explain.     Equations? 

(e)  Add   ammonium   sulphide   to   a   solution   of   an   aluminum   salt. 
Result?    Filter,  and  wash  the  precipitate  until  it  is  free  from  the  pre- 
cipitant, i.e.,  is  odorless.     Remove  the  precipitate  to  a  test  tube,  and 
add  hydrochloric  acid  (?).     Is  hydrogen  sulphide  evolved?     Was  the 
precipitate  a  sulphide? 

The  hydroxide  and  not  the  sulphide  was  formed  in  the  first  reaction 
as  the  sulphide  cannot  exist  in  the  presence  of  water.  Explain.  Express 
the  precipitation  of  aluminum  hydroxide  by  ammonium  sulphide  by 
ionic  equations. 

(/)  Mordants,  (i)  To  i  cm.3  of  a  cochineal  solution  add  5  or  6  cm.3 
of  a  solution  of  aluminum  sulphate,  then  add  ammonium  hydroxide. 
Shake  vigorously  and  filter.  Locate  the  coloring  matter  (?).  Does 
water  remove  it  easily? 

(2)  Immerse  a  strip  of  white  muslin  in  a  strong  solution  of  aluminum 
sulphate.  When  the  muslin  becomes  saturated,  transfer  it  to  a  hot 
solution  of  cochineal  which  has  been  made  strongly  alkaline  with  am- 
monium hydroxide.  What  are  mordants?  State  the  purpose  of  the 
use  of  the  ammonium  hydroxide  in  conjunction  with  the  aluminum 
sulphate. 


332  EXPERIMENTAL  CHEMISTRY. 

THALLIUM,  Tl. 

At.  Wt.  214.1     Sp.  Gr.  U.S. 
Experiment  I.  —  Properties  of  Thallium. 

Examine  a  specimen  of  the  metal  (?).  It  tarnishes  on  exposure  to 
air  with  the  formation  of  black  thallous  oxide.  Has  the  specimen  under- 
gone oxidation  ? 

In  the  univalent  condition  its  properties  suggest  those  of  sodium  and 
silver,  while  trivalent  thallium  in  its  compounds  resembles  aluminum; 
for  example,  thallic  salts  are  hydrolyzed  by  water. 

Experiment  II.  —  Salts  of  Thallium. 

Examine  those  compounds  of  the  metal  which  are  found  on  the  end 
shelf.  Observe  the  name,  formula  and  color  of  each  salt.  How  many 
well-defined  series  of  thallium  compounds  do  you  find?  What  is  the 
valence  of  thallium  in  each  series?  Tabulate  data. 

PROBLEMS. 

i.  Calculate  the  amount  of  alumnia,  A12O3,  which  can  be  prepared 
from  1000  grams  of  bauxite,  A12O(OH)4. 


A12O(OH)4  +  Na^Og  —  2NaAlO2  +  CO2  +  2H2O, 
2NaAlO2  +  CO2  +  3H2O  ->  N2CO3  +  2A1(OH)3. 
2A1(OH)3  ->  A12O3  +  3H2O. 


CHAPTER  XXIX. 
THE    ELEMENTS    OF    GROUP    IV. 


Family  M. 
(Titanium,      Ti. 
(Zirconium,     Zr. 
(Cerium, 
(Thorium, 


48.1 
90.6 

Ce.    140.25 
Th.  232.5 


Family  m. 
Carbon,  C. 

Silicon,  Si. 

(Germanium,   Gr. 
Tin,  Sn. 

Lead,  Pb. 


12  .OO 

28.4 

72.5 
II9.O 
206.9 


The  four  rare  elements,  titanium,  zirconium,  cerium  and  thorium, 
compose  Family  M  of  Group  IV. 

Titanium  as  the  constituent  of  several  rare  minerals  occurs  chiefly  in 
the  form  of  the  crystalline  titanium  oxide,  of  which  there  are  several 
native  varieties,  known  as  rutile,  anatase  and  brookite.  It  is  exceedingly 
difficult  to  prepare  the  metal  in  the  pure  form,  owing  to  the  fact  that  it 
combines  readily  with  nitrogen,  forming  a  nitride.  Zirconium  is  found 
in  the  rare  silicate,  ZrSiO4,  in  the  mineral  zircon.  This  element  resembles 
boron  in  that  it  occurs  in  several  allotropic  varieties:  an  amorphous  form, 
a  black  powder,  a  graphitic  form,  and  the  crystalline  form,  hard  steel- 
gray  lamella.  The  amorphous  form  burns  in  the  air  when  heated 
gently,  but  the  crystalline  varieties  require  the  high  temperature  of  the 
oxyhydrogen  flame  for  its  ignition.  Zirconium  dioxide  (zirconia), 
ZrO2,  which  emits  a  white  and  brilliant  light  when  heated,  is  used  as  a 
constituent  of  the  mantles  of  the  Welsbach  lights.  Cerium  occurs  chiefly 
in  the  two  rare  minerals,  cerite  and  orthite.  The  metal  as  obtained  by 
electrolysis  is  steel-gray,  lustrous  and  malleable.  It  is  stable  in  the 
air  at  ordinary  temperatures,  but  possesses  the  peculiar  property  of 
undergoing  oxidation  with  incandescence  when  heated.  In  virtue  of 
this  latter  property,  cerium  has  recently  become  of  much  technical 
importance  in  the  manufacture  of  the  Welsbach  mantle.  It  has  been 
suggested*  that  "  the  effectiveness  of  cerium  dioxide,  CeO2,  as  a  con- 
stituent of  the  Welsbach  mantle  in  intensifying  the  light  emitted  may 
depend  upon  the  facility  with  which  it  is  reduced  and  reoxidized,  thereby 
acting  as  a  catalyzer  to  cause  rapid  combustion  and  high  temperature 
in  the  flame."  Thorium,  which  until  comparatively  recently  was  obtained 
almost  solely  from  thorite  (a  silicate),  is  now  obtained  from  the  mineral 
monazite  which  is  found  in  abundance  in  North  Carolina.  Its  metallic 
properties  are  very  similar  to  those  of  cerium.  The  dioxide  of  thorium 
is  the  chief  constituent  of  the  Welsbach  mantles,  which  are  usually  pre- 
pared by  moistening  a  cotton  web  with  a  solution  containing  thorium 

*  Gooch  and  Walker — Outlines  of  Inorganic  Chemistry. 
333 


334  EXPERIMENTAL  CHEMISTRY. 

nitrate  and  one  per  cent,  of  cerium  nitrate,  then  drying  and  heating  it. 
The  web  (organic  material)  burns  and  the  thoria  and  ceria  remain  as  a 
white  coherent  frame-work.  The  mantle  is  made  ready  for  use  by 
heating  it  in  a  Bunsen  fla.me  produced  under  pressure.  The  material 
of  the  mantle  contracts  appreciably  as  the  result  of  this  last  operation. 

Although  the  exact  role  of  the  ceria  is  not  known,  yet  it  is  a  fact  that 
pure  thoria  makes  a  poorly  luminous  mantle.  Ostwald*  says:  "The 
cause  of  this  influence  (referring  to  the  luminous  properties  of  the  mantle 
when  small  quantities  of  certain  other  substances  are  mixed  with  the 
thoria)  has  not  yet  been  established  quite  free  from  doubt,  but  the  most 
probable  view  is  that  the  addition  (of  the  ceria)  effects  a  catalytic  acceler- 
ation of  the  combustion  of  the  mixture  of  coal-gas  and  air  in  direct  contact 
with  the  skeleton  of  thoria.  Besides  this,  the  optical  properties  of  thoria 
appear  to  be  of  importance." 

The  compounds  of  thorium  are  radio-active;  i.e.,  they  possess  the 
power  of  continually  emitting  certain  "influences,"  sometimes  called 
the  "Becquerel  rays"  and  "emanations"  which  possess  the  property 
of  effecting  a  change  in  the  electrical  properties  of  the  air.  Certain  of 
these  "influences"  are  able  to  penetrate  solid  substances  and  affect  a 
photographic  plate.  The  student  is  referred  to  reference  works  on  this 
subject  of  radio  activity,  f 

All  the  metals  of  this  family  are  quadrivalent. 

The  elements  of  Family  m.  have  been  divided  into  two  series  (see 
Chapter  XX);  a  primary  series  embracing  carbon  and  silicon,  and  a 
secondary  series  composed  of  the  elements,  germanium,  tin  and  lead. 
The  primary  series  has  been  considered  previously. 

The  elements  of  the  secondary  series  are  silver-wrhite,  lustrous,  mal- 
leable metals  which  are  practically  unaffected  by  air  or  water.  The 
metals  are  fusible  and  volatile  at  high  temperatures.  They  possess 
high  specific  gravities  which  increase  as  the  atomic  weights  increase. 

The  rare  element  germanium  (Mendeleeff's  eka-silicon)  forms  a 
sort  of  link  between  the  elements  of  the  primary  series  and  those  of 
the  secondary  series.  It  is  both  metallic  and  non-metallic;  its  oxide 
combines  with  acids,  but  it  also  unites  with  alkaline  hydroxides.  Both 
tin  and  lead  which  resemble  each  other,  especially  in  their  physical 
properties,  show  similar  basic  and  acidic  relations  in  their  respective 
compounds;  i.e.,  their  oxides  and  hydroxides  combine  with  strong  acids 
to  form  salts  in  which  tin  and  lead  are  positive  ions,  while  with  the 
caustic  alkalies  they  form  compounds,  known  respectively  as  stannates 
of  the  type,  NaSnO2,  and  plumbites  of  the  type,  Na-jPbO-j,  in  which  the 
metals  are  in  the  negative  ions.  Each  element  of  the  series  forms  com- 
pounds of  the  types  stannous  chloride  (SnCl2)  and  stannic  thloride  (SnCl4), 
in  which  the  metals  function,  respectively,  as  bivalent  and  quadrivalent 

*  Prin.  of  Inorg.  Chem. 

t  Rutherford — Radioactivity;  Arrhenius — Theories  of  Chemistry;  Fournier — The 
Electron  Theory. 


THE    ELEMENTS    OF    GROUP    IV.  335 

elements.  Lead,  however,  manifests  a  greater  tendency  to  react  at  the 
state  of  lower  valency.  This  is  not  an  exception,  but  rather  a  rule  that 
the  heavier  metals  of  a  group  tend  to  react  at  the  lower  valency. 

The  compounds  of  these  three  elements,  many  of  which  are  insoluble, 
are  but  slightly  hydrolyzed  by  water  and  generally  stable. 

The  chief  ore  of  tin  is  tin-stone,  or  cassiterite,  SnO2,  the  greater  portion 
of  which  (fully  80  per  cent.)  is  supplied  by  Cornwall  and  the  East  Indies. 
The  metallurgy  of  tin  consists  of  four  processes,  namely,  (i)  crushing 
or  pulverizing  the  ore,  (2)  calcining,  (3)  washing,  (4)  smelting  or  re- 
ducing. After  the  ore  has  been  crushed  finely  and  washed  to  free  it 
from  earthy  matter,  it  is  calcined  in  a  reverberatory  furnace.  The 
latter  operation  is  for  the  purpose  of  oxidizing  the  sulphides  of  iron  and 
copper  and  to  drive  off  the  arsenic.  The  sulphur  and  arsenic  are  led 
into  condensing  flues  where  the  arsenic  deposits  are  collected.  The 
calcined  ore  is  now  washed  to  eliminate  the  oxide  of  iron  and  the  sulphate 
of  copper,  and  then  reduced  with  powdered  anthracite  coal  in  a  rever- 
beratory furnace.  The  tin  obtained  by  the  foregoing  process  is  further 
purified  by  remelting  at  a  gentle  heat,  the  pure  and  more  readily  fusible 
tin  being  allowed  to  flow  away  from  the  residue  and  alloys  of  other 
metals. 

The  crystalline  character  of  the  metal  may  be  observed  by  pouring 
warm  dilute  aqua  regia  over  the  surface  of  a  piece  of  block-tin  or  a  sheet 
of  tinned  iron.  The  surface  of  the  metal  after  such  treatment  exhibits 
a  beautiful  crystalline  appearance.  Again,  when  a  bar  of  tin  is  bent 
a  peculiar  crackling  sound  (tin  cry)  is  produced.  It  has  been  observed 
also  that  the  metal  becomes  perceptibly  hot  at  the  place  of  flexure.  It 
has  been  suggested  that  the  cause  of  these  phenomena  may  be  due  to  the 
friction  of  the  crystals  upon  one  another.  When  ordinary  tin,  which 
has  a  specific  gravity  of  7.3.  is  exposed  to  the  prolonged  influence  of  low 
temperature,  it  changes  to  a  gray  pulverulent  variety  of  specific  gravity 
5.8.  The  transition  temperature  is  20°,  and  ordinary  tin  is  in  a  meta- 
stable  condition  below  this  temperature.  Tin  is  not  tarnished  by  air 
or  water  at  ordinary  temperatures,  but  when  heated  above  the  melting- 
point  it  burns  with  a  brilliant  white  light,  forming  white  clouds  of  stannic, 
SnO2.  At  red  heat  tin  decomposes  steam  with  a  liberation  of  hydrogen. 
Hydrochloric  acid  and  tin  interact  with  the  evolution  of  hydrogen  and 
the  formation  of  stannous  chloride;  when  strong  sulphuric  acid  is  heated 
writh  tin,  stannous  sulphates  and  sulphur  dioxide  are  formed;  cold 
dilute  nitric  acid  yields  stannous  nitrate  and  ammonia,  while  concen- 
trated nitric  acid  gives  stannic  nitrate,  which  is  hydrolyzed  by  water  with 
the  formation  of  metastannic  acid  (H2SnO3)5.  The  caustic  alkalies 
attack  tin,  giving  hydrogen  and  a  metastannate,  such  as  Na2SnO3.  Tin 
forms  alloys  with  lead,  copper,  antimony,  bismuth  and  mercury.  Among 
the  more  familiar  of  these  alloys  are  bronze,  sojt  solder  (50  per  cent, 
lead),  pewter  (25  per  cent,  lead),  Britannia  metal  (10  per  cent,  antimony 
and  a  small  quantity  of  copper),  and  fusible  alloys  (bismuth,  tin,  cadmium 
and  sometimes  lead).  Tin  is  also  used  as  a  protective  covering  on  other 


33  EXPERIMENTAL  CHEMISTRY. 

metals  on  account  of  the  difficulty  with  which  it  is  attacked  by  many 
corroding  substances.  Tin-plate  is  made  by  dipping  sheet  iron  into 
molten  tin.  Vessels  made  of  copper  are  also  frequently  covered.  Or- 
dinary brass  pins  are  made  of  brass  wire  coated  with  tin.  The  metal 
finds  many  other  uses  in  the  arts. 

Lead  is  obtained  almost  wholly  from  the  ore,  galena  PbS.  The 
metallurgical  processes  by  which  lead  is  obtained  from  its  ores  are 
similar  to  those  described  for  the  reduction  of  sulphides.  The  metal 
is  refined  by  electrolytic  methods.  It  is  a  grayish-white  metal,  soft  and 
tough.  It  melts  at  about  330°,  and  vaporizes  at  1700°.  Under  suitable 
conditions  it  crystallizes,  the  crystals  having  the  octahedral  form.  When 
warm,  it  may  be  formed  into  pipes  by  hydraulic  pressure.  It  is  a  poor 
conductor  of  electricity.  Lead  is  oxidized  but  superficially  by  the  air, 
becoming  covered  with  a  film  of  a  dark-colored  oxide,  which  is  probably 
the  suboxide  Pb2O,  the  composition  of  the  final  covering  being  that  of  a 
basic  carbonate;  when  heated,  the  metal  passes  through  several  stages  of 
oxidation  with  the  formation  of  no  less  than  five  distinct  oxides.  Pure 
water  does  not  act  upon  lead,  but  hard  water  covers  it  with  a  coating 
composed  largely  of  the  sulphate  and  the  carbonate.  As  these  salts 
are  insoluble,  they  protect  the  metal  and  prevent  contamination  of  the 
water  with  poisonous  lead  compounds.  Water  holding  air  in  solution 
attacks  lead,  forming  the  slightly  soluble  hydroxide  and  the  carbonate. 
The  latter  is  appreciably  soluble  in  water  containing  carbon  dioxide. 
The  use  of  lead  pipes  for  conducting  water  may  become  a  source  of 
danger,  as  rain  water  especially,  owing  to  the  presence  of  oxygen  and 
carbon  dioxide,  is  likely  to  exert  a  solvent  action  upon  the  pipes.  The 
lead  compounds  taken  continuously  into  the  system  in  small  quantities 
act  as  a  cumulative  poison.  Hydrochloric  acid  acts  upon  the  metal 
slowly  with  an  evolution  of  hydrogen.  Concentrated  sulphuric  acid 
has  but  little  affect  upon  it;  dilute  sulphuric  acid  slowly  interacts  with  it, 
forming  the  insoluble  lead  sulphate  and  sulphur  dioxide.  Nitric  acid 
attacks  it,  giving  lead  nitrate  and  oxides  of  nitrogen.  The  action  of  the 
caustic  alkalies  upon  lead  has  been  mentioned  previously. 

The  uses  of  lead  are  numerous.  It  forms  useful  alloys  with  tin, 
antimony,  bismuth,  copper  and  zinc.  Type-metal  usually  contains  lead 
and  antimony  and  sometimes  tin.  On  account  of  the  resistance  of  lead 
to  the  action  of  air  and  water  and  many  other  substances,  it  is  employed 
in  making  various  kinds  of  vessels  and  lead  pipes  for  carrying  water. 
Shot-metal  is  an  alloy  composed  of  about  99.5  per  cent,  of  lead  and  0.5 
per  cent,  of  arsenic.  "  In  the  process  of  making  shot  the  lead  is  melted 
in  a  cast-iron  pan,  and  after  the  addition  of  a  sufficient  amount  of  lead 
arsenide  to  form  an  alloy  containing  from  0.8  to  0.9  per  cent,  of  arsenic, 
the  molten  alloy  is  poured  into  a  perforated  iron  basin  at  the  top  of  a 
high  tower,  and  allowed  to  fall  into  a  dilute  solution  of  sodium  sulphide. 
The  presence  of  arsenic  makes  the  drops  of  molten  alloy  very  fluid,  so 
that  they  assume  a  spherical  form  in  their  passage  to  the  bottom  of  the 
tower.  The  effect  of  the  sodium  sulphide  is  to  coat  the  shot  with  a  thin 


THE    ELEMENTS    OF    GEOUP    IV. 


337 


layer  of  lead  sulphide  which  prevents  superficial  oxidation  when  the  shot 
are  removed  from  the  water.  Shot  are  sometimes  made  by  allowing  the 
molten  metal  to  fall  through  an  ascending  air  current  or  by  pouring 
it  in  a  thin  stream  upon  a  rapidly  revolving  disk,  when  the  centrifugal 
action  divides  the  metal  into  drops,  which  are  thrown  against  a  surround- 
ing screen."* 

The  following  table  gives  a  general  view  of  the  physical  and  chemical 
properties  of  the  metals  of  this  series: 


Atomic  weight, 
Specific  gravity, 
Melting-point, 
Volatilizes, 

Oxides, 


Chlorides, 

Hydroxides, 
Sulphides, 


Germanium. 

72-5 

5-47 
900° 


Tin. 

119.0 


231 


IS00°(?) 


Lead. 
206.9 

n-37 
330° 
1700° 


GeO;  GeO2,    SnO;  SnO2,     Pb2O;  PbO;  Pb2O3; 

PbO2;  Pb3O4. 

GeCl2;GeCl4,     SnCl2;  SnCl4,     PbCl2;  PbCl4. 

Ge(OH).,;  Ge(OH4,  Sn(OH)2,Sn(OH)4,  Pb(OH)2. 
GeS;  GeS2          SnS;  SnS2,          Pb2S;  PbS. 


GERMANIUM,    GC. 

At.  Wt.  72.5         Sp.  Gr.  7.3. 
(See  lecture  notes  and  text-book.) 


At.  Wt.  119.0         Sp.  Gr.  7.3. 
Experiment  I. — Properties  of  Tin. 

(a)  Examine  a  specimen  of  this  metal  and  record  its  most  obvious 
physical  properties.     What  is  its  melting  point  ?     Is  it  hard  or  soft  ? 

(b)  (Quant.)     Determine  the  specific  heat  of  tin.     Secure  the  directions 
for  procedure  from  the  assistant.     Calculate  the  atomic  weight  of  tin, 
using  0.054  as  its  specific  heat. 

(c)  Does  tin  tarnish  readily  wrhen  exposed  to  the  action  of  the  air? 
Enumerate  some  of  the  uses  of  tin.     Name  its  chief  ores. 


Experiment  II.— (Quant.) 
Tin. 


Determination  of  the  Equivalent  Weight  of 


In  a  crucible  which  has  been  previously  desiccated  and  weighed, 
place  an  accurately  weighed  quantity  (about  0.5  gram.)  of  pure  granulated 
tin,  and  cover  the  metal  with  about  10  cm.3  of  concentrated  nitric  acid. 


*  Gooch  and  Walker. — Outlines  of  Inorganic  Chemistry. 


22 


338  EXPERIMENTAL  CHEMISTRY. 

Carefully  apply  heat  (use  a  hot  iron  plate)  to  the  crucible.  Identify  the 
reddish-brown  fumes.  When  all  of  the  tin  has  dissolved  and  the  nitric 
acid  has  been  entirely  expelled,  place  the  crucible  on  a  pipe-stem  triangle 
and  heat  it  with  a  Bunsen  burner.  The  white  product  is  stannic  oxide, 
SnO2.  Equation?  Cool,  and  weigh  the  crucible  and  contents.  Cal- 
culate the  chemical  equivalent  of  tin.  If  the  valence  of  tin  in  stannic 
oxide  is  four,  what  is  its  atomic  weight  ? 

Experiment  III. — Formation  of  Halides  of  Tin. 

(a)  Stannous  chloride.     Treat  several  small  pieces  of  granulated  tin 
with  10  cm.3  of  concentrated  hydrochloric  acid  in  a  test  tube,  and  warm 
gently  to  start  the  action.     After  the  action  has  continued  for  some  time, 
pour  off  the  liquid  into  another  test  tube  and  reserve  it  for  use  in  those 
experiments  in  which  stannous  chloride,  SnCl.2,  is  required.     Equation  ? 

(b)  Stannic  chloride.     Add  3  cm. 3  of  concentrated  nitric  acid  and  i 
cm.3  of  hydrochloric  acid  to  5  cm*3  of  stannous  chloride,  and  heat  gently. 
The  solution  contains  stannic  chloride,   SnCl^.     Dilute  with   5  cm.3  of 
water.     Equations  ? 

Experiment  IV. — Precipitation  of  Stannous  Hydroxide.  Sodium  Stan- 
nite. 

To  5  cm.3  of  a  dilute  solution  of  stannous  chloride  add  carefully  a 
sodium  hyrdoxide  solution  until  precipitation  is  complete.  Equation  ? 

Divide  the  precipitate  into  two  parts.  To  one  portion  add  an  excess 
of  the  sodium  hydroxide  solution  ( ?)  and  to  the  other  portion  add 
hydrochloric  acid  (?).  Equations? 

Suggest  a  method  for  the  preparation  of  stannous  oxide,  SnO.  Equa- 
tion? 

Experiment  V. — Precipitation  of  Stannic  Hydroxide  (a — Stannic  Acid). 
Sodium  Stannate. 

Repeat  Experiment  IV  substituting  stannic  chloride  for  stannous 
chloride. 

What  product  is  obtained  by  heating  stannic  hydroxide  in  a  crucible  ? 
Equation  ? 

Epperiment  VI. — Analytical  Reactions. 

(a)  Heat  a  small  piece  of  tin  on  charcoal  before  the  blow-pipe  (?). 
Allow  a  drop  of  cobalt  nitrate  to  fall  upon  the  incrustation,  then  heat 
again  (?). 

(b)  Place  a  piece  of  zinc   or  suspend  a  strip  of  it  in  a  solution  of 
stannous  chloride  for  ten  to  fifteen  minutes.     Result  ?     Equation  ? 

(c)  Add  i  cm.3  of  a  solution  of  mercuric  chloride  to  4  or  5  cm.3  of 
stannous    chloride     solution.     Result?     The     solution    now    contains 
stannic  chloride,  SnCl4.     Equation?     Warm  the  mixture  and  note  the 
changes  ( ?).     Equation  ?     Does  the  stannous  chloride  act  as  a  reducing 
or  an  oxidizing  agent?     Indicate  all  of  the  above  changes  by  "ionic" 
equations.     Repeat  the  foregoing  reaction,  using  stannic  chloride  instead 


THE    ELEMENTS    OF    GROUP   IV.  33 Q 

of  stannous  chloride  (?).     How  can  you  distinguish  between  stannous 
and  stannic  compounds? 

(d)  Add  a  few  drops  of  stannic  chloride  to  5  cm.3  of  dilute  hydro- 
chloric  acid  in   a   test  tube.     Now   add   several  pieces  of   magnesium 
ribbon  to  the  acid  solution.     What  is  evolved?     Is  it  an  oxidizing  or 
reducing  agent  when  in  the  nascent  condition?     When  the  magnesium 
has  dissolved,  add  a  small  quantity  of  the  mercuric  chloride  solution. 
Result  ?     Conclusions  ?     Equations  ? 

(e)  Pass  hydrogen  sulphide  into  5  cm.3  of  absolution  of  stannous  chloride 
containing    i    cm.3    of   dilute    hydrochloric    acid.     Result?     Equation? 
Is  the  reaction  easily  reversible  ?     Give  reasons  for  your  answer.     Filter 
and  wash  the  precipitate.     Place  the  latter  in  an  evaporating  dish,  add 
10  cm.3  of  ammonium   polysulphide   and  warm.     Result?     Equation? 
Now  add  hydrochloric  acid  to  the  solution  of  ammonium  sulphostannate, 
(NH4)2SnS3.     Results?     What  gas  was  evolved?     Equation? 

Repeat  the  preceding  reactions,  using  stannic  chloride  (?). 
(/)  There  are  how  many  ionic  forms  of  tin  ?     What  is  the  color  of  the 
tin  ions  ?     Give  the  valencies  of  the  respective  ions. 

LEAD,    Pb. 

At.  Wt.  206.9         Sp.  Gr.  11.37. 
Experiment  I. — Properties  of  Lead. 

(a)  File  or  scrape  off  the  coating  from  a  piece  of  lead.     Is  the  metal 
hard  or  soft?     Color?     Does  lead  tarnish  (oxidize)  readily  in  the  air? 
Try  to  mark  on  paper  with  lead  ( ?). 

(b)  Solution  tension.     Dissolve  about  i  gram  of  lead  acetate,  Pb(C2H3- 
O2)2,  in  20  cm.3  of  water.     Place  a  strip  of  sheet  zinc  or  several  pieces  of 
granulated  zinc  in  the  solution,  and  set  aside  for  an  hour.     Result? 
Ionic  equation  ?     Remove  the  film  of  lead  from  the  zinc,  thoroughly 
wash  the  former  and  reserve  it  for  (c). 

(c)  Action  of  air  and  water  on  lead.     Spread  the  finely  divided  lead 
from  (b)  on  a  glass  plate,  moisten  with  a  very  little  water,  and  expose  to 
the  action  of  the  air  for  an  hour.     Test  the  water  with  litmus  papers  (  ?). 
Conclusions?     The  presence  of  lead  in  the  water  may  be  detected  by 
shaking  a  portion  of  the  lead  with  a  small  volume  of  water  and  passing 
hydrogen  sulphide  into  the  filtrate.     A  black  precipitate  (lead  sulphide) 
indicates  the  presence  of  lead.     Equations  ? 

(d)  Name  the  chief  ore  of  lead.     What  are  some  of  the  uses  of  lead? 
What  is  its  melting  point  ?     What  is  pewter  ? 

Experiment  II. — (L.  T.)  Precipitation  of  Lead  from  its  Salts  by  Other 
Metals.  "Lead  Tree." 

Suspend  a  bar  of  zinc  in  a  solution  of  lead  acetate  or  nitrate.  Allow 
this  to  stand  for  several  days.  The  lead  will  be  gradually  thrown  out 
of  solution  (why?)  and  deposited  upon  the  zinc  in  arborescent  forms, 


34°  EXPERIMENTAL  CHEMISTRY. 

known  as  the  "lead  tree."  Test  the  solution  for  the  presence  of  zinc, 
as  follows:  Add  sufficient  sulphuric  acid  to  precipitate  any  lead  remain- 
ing in  solution,  filter,  and  to  the  filtrate  add  ammonia  in  excess  and 
ammonium  sulphide;  the  zinc  will  be  precipitated  as  sulphide.  Ionic 
equations  ? 

Experiment  III. — Salts  of  Lead. 

Examine  a  number  of  the  salts  of  lead.  Give  color  and  formula  of 
each.  What  is  the  valency  of  lead  ? 

Note. — Lead  forms  five  oxides;  give  the  name  and  formula  of  each. 

Experiment  IV. — Precipitation  of  Lead  Hydroxide.     Sodium  Plumbate. 

(a)  To  a  solution  of  lead  nitrate  add  at  first  slowly  and  then  in  excess 
a  sodium  hydroxide  solution.     Describe  all  of  the  changes  that  occur. 
Ionic  equations?     Does  lead  exhibit  the  properties  of  both  a  metal  and 
non-metal  ?     Give  reasons  for  your  answer. 

(b)  Repeat  (a)  using  ammonium  hydroxide. 

Experiment  V. — Preparation  of  Lead  Salts. 

(a)  Lead  nitrate  from  lead  and  nitric  acid.     Treat  about  a  gram  of 
lead  in  an  evaporating  dish  with  20  cm. 3  of  a  mixture  (i  to  2)  of  nitric 
acid  and  water.     Place  the  dish  on  a  wire  gauze  and  heat  gently  until  the 
metal  dissolves.     Note  the  accompanying  phenomena.     Set  the  solu- 
tion aside  to  crystallize  by  spontaneous  evaporation.     Result?     Equa- 
tion? 

(b)  Lead  nitrate  from  lead  monoxide  (litharge)  and  nitric  acid.     Dis- 
solve 3  to  5  grams  of  lead  monoxide  in  dilute  nitric  acid,  filter,  evaporate 
the  filtrate  to  the  crystallizing  point,  and  cool.     Result  ?     Equation  ? 

(c)  Lead  nitrate  and  lead  dioxide  from  red-lead  oxide  (minium)  and 
nitric  acid.     Treat  about  a  gram  of  minium  with  dilute  nitric   acid, 
warm  gently,  and  when  the  red  color  of  the  minium  has  changed  to  a 
brown,  dilute  with  water  and  filter.     Wash  the  residue  ( ?)  and  ignite 
it.     Which  of  the  oxides  of  lead  does  it  resemble  most  in  appearance? 
Prove  the  presence  of  lead  in  the  filtrate  by  adding  slowly  an  excess  of  a 
solution    of    sodium    hydroxide    (?).     Equations?     This    behavior    of 
minium  suggests  what  theory,  as  to  its  constitution  ? 

(d)  Lead  acetate  from  lead  oxide  and  acetic  acid.     To  5  grams  of  lead 
monoxide  add  10  cm.3  of  acetic  acid,  then  warm  gently  to  increase  the 
speed  of  the  reaction.     If  the  solution  is  not  clear,  filter,  and  evaporate 
to  the  crystallizing  point,   exercising  care  to  avoid  charring  the  salt. 
Equations?     What  is  "sugar  of  lead?" 

Experiment  VI. — Precipitation  of  Lead  Carbonate.  Basic  Lead  Car- 
bonate ;  its  Decomposition  by  Heating. 

(a)  When  a  solution  of  ammonium  carbonate  is  added  to  a  solution  of 
lead  nitrate,  the  normal  carbonate  is  formed.  (Equations  ?)  If  any 
other  alkaline  carbonate  is  used,  a  basic  lead  carbonate  is  formed. 


THE    ELEMENTS    OF    GROUP    IV.  341 

(b)  To  10  cm.3  of  a  solution  of  lead  nitrate  add  an  excess  of  sodium 
carbonate  solution  (?).  Filter,  wash  and  dry  the  precipitate,  separate 
it  from  the  filter  paper,  and  ignite  the  white  powder  in  a  porcelain  crucible. 
Avoid  heating  the  crucible  above  dull  redness.  Observe  the  color  of  the 
residue  (?)  when  hot  and  cold.  Equations?  What  is  "white  lead"? 
Its  use? 

Experiment  VII. — Chemical  Principles  Involved  in  the  Manufacture 
of  "  White  Lead." 

Add  a  little  lead  monoxide  to  10  cm.3  of  a  solution  of  lead  acetate; 
thoroughly  shake  the  mixture;  pass  carbon  dioxide  (generator  on  end 
shelf)  through  it  until  it  is  white;  filter.  The  precipitate  is  used  as  a 
pigment  under  the  name  of  "white  lead."  Equations?  What  is  the 
effect  of  hydrogen  sulphide  upon  white  lead  ?  Equation  ? 

Experiment  VIII. — Analytical  Reactions. 

(a)  Heat  a  small  piece  of  lead  on  charcoal  in  the  oxidizing  flame 
(blow-pipe).     Result?     Equation? 

(b)  Test  a  solution  of  lead  nitrate  with  litmus  paper  (?).     Interpret 
the  result. 

(c)  Halides  of  lead.     To  2  cm.3  of  a  solution  of  lead  nitrate  add  dilute 
hydrochloric    acid   until   precipitation   is   complete?     Filter,    wash    the 
precipitate,  and  heat  it  with  the  smallest  volume  of  water  which  will 
dissolve  it.     Cool  the  solution.     Result  ?     Equations  ? 

In  what  two  respects  does  lead  chloride  differ  from  silver  chloride? 
Name  the  insoluble  chlorides. 

Treat  a  cm.3  of  lead  nitrate  solution  with  a  solution  of  potassium  or 
sodium  bromide.  Results  ?  Equation  ?  Is  the  precipitate  soluble  in 
hot  water  ?  Compare  its  solubility  with  that  of  lead  chloride. 

Add  a  solution  of  potassium  iodide  to  2  cm.3  of  lead  nitrate  solution. 
Results  ?  Equation  ? 

How  does  the  solubility  of  the  precipitate  compare  with  the  solubility 
of  lead  bromide? 

(d)  To  2  cm.3  of  lead  nitrate  solution  add  dilute  sulphuric  acid.     Re- 
sult?    Equation.     What  sulphates  are  insoluble?     Give  the  formula  of 
each. 

(e)  Allow    hydrogen    sulphide    (generator  on    end  shelf)    to    bubble 
slowly  through  5  cm.3  of  a  solution  of  lead  nitrate  to  which  has  been 
added  four  or  five  drops  of  dilute  nitric   acid.     Result?     Equation? 
Is  the  action  reversed  easily  ?     Give  reasons  for  your  answer. 

(/)  Add  a  dilute  solution  of  potassium  chromate  to  3  cm.3  of  lead 
nitrate  solution.  Result?  Equation?  Is  the  precipitate  soluble  in  a 
strong  solution  of  sodium  hydroxide?  Equation?  What  is  "chrome 
yellow"  ?  Its  uses? 

(g)  What  is  the  color  of  the  lead  ion  ?  Does  it  ever  form  a  part  of  the 
anion  ?  If  so,  give  an  example. 


342  EXPERIMENTAL  CHEMISTRY. 

PROBLEMS. 

1.  Calculate  the  weight  of  tin  in  i  kg.  of  cassiterite,  SnO2. 

2.  What  per  cent,  of  lead  is  contained  in  galena,  PbS  ? 

3.  If  50  grams  of  tin  yields  63.55  grams  of  stannic  oxide,  what  is 
the  atomic  weight  of  tin? 

4.  An  analysis    showed  that    i    gram  of    lead  monoxide  contained 
0.0717  gram  of  oxygen.     Calculate  the  atomic  weight  of  lead. 

5.  The  per  cent,   of   lead  in   lead  chloride  is  74.4    and  the  specific 
heat  of  the  metal  is  0.031.     Calculate  the  atomic  weight  of  lead. 


CHAPTER   XXX. 
ELEMENTS  OF  GROUP  V. 

Family  M. 

(Vanadium,  V.  51.2) 
(Columbium,Cb.  94.  ) 
(Tantalum,  Ta.  181.  ) 

The  elements  of  this  family  are  lustrous  gray  solids  which  are  very  rare 
and  difficult  to  isolate.  Vanadium  is  the  least  uncommon.  The  mem- 
bers of  the  family  are  closely  related  to  one  another. 

Vanadium,  which  was  first  isolated  by  Roscoe  in  1867,  is  found  in  the 
complex  mineral,  vanadinite,  Pb4(PbCl)  (VO4)3.  It  is  slowly  acted  upon 
by  air  at  ordinary  temperatures,  but  when  heated  it  burns  brilliantly, 
forming  the  reddish-brown  vanadium  pentoxide,  V2O5.  This  oxide  in- 
teracts with  bases  giving  vanadaies.  The  metal  possesses  very  feeble 
base-forming  properties.  Vanadium  also  combines  with  nitrogen  at  red 
heat  to  form  the  yellowish-red  vanadium  nitride,  VN.  The  following 
are  among  the  more  important  of  its  compounds:  V2O,  V2O2,  V2O3, 
VO2,  V2O5,  VC12,  VC13,  VC14,  VOC13,  VOC15,  V2S,  V2S3,  and  H3VO4. 
The  element  is  frequently  prepared  by  heating  VC12  in  a  stream  of 
hydrogen. 

Columbium  (niobium)  and  tantalum  are  found  in  the  rare  minerals 
columbite,  (Mn,Fe)  (Cb,Ta)2O6,  and  tantalite,  (Fe,Mn)Ta2O6,  respec- 
tively. These  two  elements  likewise  possess  feebly  base-forming  properties. 
The  chief  compounds  are  the  columbates  and  tantalates.  Other  com- 
pounds are,  Cb2O2,  Cb2O4,  Cb2O5,  CbOCl3,  CbCl3,  CbCl5,  H3CbO4, 
Ta2O4,  Ta2O5,  TaCl5  and  H3TaO4. 


343 


CHAPTER  XXXI. 
ELEMENTS  OF  GROUP  VI. 

Family  M. 

Chromium,        Cr.     52.1 
(Molybdenum,  Mo.    96.0)- 
(Tungsten,         W.    184.0) 
(Uranium,          U.     238.5) 

The  elements  of  this  group  possess  many  properties  in  common.  They 
are  hard  metals  possessing  a  metallic  luster  and  a  high  specific  gravity. 
They  are  attacked  by  acids,  but  are  not  acted  upon  by  air  and  water  at 
ordinary  temperatures.  They  are  readily  oxidized  or  redaced  and  their 
corresponding  compounds  are  structurally  and  chemically  similar.  Their 
higher  oxides  are  acid  anhydrides,  while  their  lower  oxides  are  basic  in 
character.  They  form  two  classes  of  salts,  the  "ous"  and  the  "ic"; 
also  two  classes  of  compounds,  "ites"  and  "ates,"  in  which  the  elements 
appear  in  the  negative  ions. 

The  maximum  valence  shown  by  the  elements  of  this  group  is  VI. 
Chromium,  however,  in  perchromic  acid,  H2Cr2O8,  reaches  a  valence  of 
VII. 

CHROMIUM,     Cr. 
At.  Wt.  52.1         Sp.   Gr.  6.92 

Chromium  does  not  occur  in  nature  in  the  uncombined  condition,  and 
its  natural  compounds  are  neither  abundant  nor  widely  distributed.  The 
chief  source  of  chromium  is  chrome  iron  ore,  or  chromate,  FeCr2O4.  It 
is  prepared  most  conveniently  by  the  "  Goldschmidt"  process.  Chro- 
mium prepared  by  this  method  is  "  passive  "  and  does  not  displace  hydrogen 
from  hydrochloric  acid  until  it  is  warmed  with  the  acid.  When  removed 
from  the  acid  and  left  in  the  air,  it  changes  slowly  into  the  inactive  form 
again.  Chromium  fuses  in  the  electric  arc,  but  not  in  the  oxyhydrogen 
flame. 

Chromium  forms  the  following  oxides  and  hydroxides:  hypothetical 
chromous  oxide,  CrO,  and  chromous  hydroxide,  Cr(OH)2,  which  are 
distinctly  basic;  chromic  oxide,  Cr2O3,  and  chromic  hydroxide,  Cr(OH)3, 
which  are  weakly  basic;  chromic  oxyhydroxide,  HCrO2,  which  is  acidic; 
chromium  trioxide,  Cr2O3,  a  water  solution  of  which  is  called  chromic  acid, 
H2CrO4.* 

It  is  obvious  that  chromium  gives  four  classes  of  compounds,  chromous 
and  chromic  salts  which  correspond  to   Cr(OH)2  and   Cr(OH)3,   and 
*  The  compound  H2CrO4  has  not  been  isolated. 
344 


ELEMENTS    OF    GROUP    VI.  345 

chromites  and  chromates,  to  HCrO2  and  H2CrO4,  respectively.  There 
is  also  another  class,  perchromates,  which  correspond  to  perchromic  acid, 
H2Cr2O8. 

The  chromous  salts  are  not  stable  in  the  air  owing  to  the  ease  with 
which  they  pass  into  the  condition  of  higher  oxidation  of  chromic  com- 
pounds. All  of  the  chromic  salts  are  hydrolyzed,  CrCl3  and  Cr2(SO4)3  to 
a  small  degree,  while  neither  the  carbonate  nor  sulphide  is  stable  in  water. 
(In  this  respect  chromic  salts  are  very  similar  to  corresponding  aluminum 
salts.)  The  perchromates  are  unstable  at  ordinary  temperatures. 

It  is  to  be  noted  that  those  chromium  ions  of  "  lower  valence,  or  with 
the  smaller  electrical  charge,  are  basic;  but  as  the  valence  increases,  or 
as  the  amount  of  electrical  energy  which  they  carry  increases,  the  basic 
property  becomes  less  and  less,  and  acidic  properties  begin  to  manifest 
themselves."* 

Experiments  I. — Properties  of  Chromium. 

Examine  the  metal  and  note  its  most  obvious  physical  properties. 
Does  it  tarnish  in  the  air? 

Experiment  II.— Compounds  of  Chromium. 

(a)  Examine  the  compounds  of  chromium  (end  shelf).     Give  the  color 
and  structural  formula  of  each  compound.     Record  the  valence  of  the 
chromium  atom  in  each  compound,  and  state  whether  it  behaves  as  an 
acid-forming  or  base-forming  element.     Tabulate  the  foregoing  data. 

(b)  What  is  the  formula,  color  and  electrical  charge  of  each  of  the 
following   ions:    the    dichromate    ion    and    the    chromic    ion.       Hint: 
Examine  solutions  of  the  salts  giving  these  ions  (end  shelf). 

Experiment  III. — Preparation  of  a  Chromate. 

Mix  5  grams  of  potassium  carbonate  with  equal  amounts  of  potassium 
nitrate  (supplies  oxygen)  and  potassium  hydroxide  in  an  iron  crucible 
(T.O.).  Heat  the  mixture  at  low  temperature  until  it  melts,  then  stir 
(use  iron  rod)  in  5  grams  of  powdered  chromite.  Heat  strongly  in  the 
flame  of  a  blast-lamp  until  further  heating  produces  no  more  change. 
Allow  the  crucible  to  cool,  then  dissolve  the  contents  in  a  little  boiling 
water.  Filter.  What  is  the  color  of  the  solution  ?  The  color  is  due  to 
the  presence  of  what  ion  ? 

4(FeO,Cr203)  +  8(K2O,CQ2)  +  yO2 ->  8(K2O,CrO3)  +  2Fe2O3  -f  8CO2. 

Experiment  IV. — Formation  of  a  Dichromate  from  a  Chromate. 

To  5  cm.3  of  potassium  chromate  solution  add  sulphuric  acid,  drop  by 
drop,  until  the  yellow  color  changes  to  orange.  The  color  is  due  to  the 
presence  of  what  ion?  Evaporate  in  small  beaker  to  crystallizing  point; 
cool  slowly.  Note  color  and  shape  of  crystals. 

*  Jones,  Elements  of  Inorganic  Chemistry 


34^  EXPERIMENTAL  CHEMISTRY. 

Interpret  the  following  equations: 

K2CrO4  +  H2CrO4  —  K2Cr2O7  +  H2O. 
2K2CrO4  +  H2SO4   —  K2Cr2O7  +  K2SO4  +  H2O. 

Experiment  V. — Formation  of  a  Chromate  from  a  Bichromate. 

To  5  cm.3  of  potassium  dichromate  solution  add  enough  sodium  hy- 
droxide solution  to  turn  the  color  yellow.  The  color  is  due  to  the  presence 
of  what  ion  ?  What  salt  has  been  formed  by  the  interaction  ?  Equation  ? 
Evaporate  until  a  crust  forms,  then  set  aside  to  crystallize.  Record  color 
of  crystals. 

Experiment  VI. — Formation  of  Chromium  Trioxide  (Chromic  Anhy- 
dride); its  Oxidizing  Power.  Chromic  Acid.* 

Prepare  about  10  cm.3  of  a  warm  saturated  aqueous  solution  of  potas- 
sium dichromate  in  a  beaker.  Filter  the  solution  if  it  is  not  clear.  Now 
add  carefully,  a  drop  at  a  time,  an  equal  volume  of  concentrated  sul- 
phuric acid.  Red  needle-shaped  crystals  of  chromium  trioxide  will 
separate.*  Equation?  When  the  mixture  has  cooled,  filter  through  a 
plug  of  glass  wool  or  asbestos  (not  filter  paper).  (Caution:  Do  not 
allow  the  crystals  to  come  in  contact  with  the  hand.}  By  means  of  a  glass 
or  porcelain  spatula  remove  a  few  of  the  crystals  and  place  them  upon  a 
piece  of  filter  paper.  Account  for  the  result.  Treat  a  few  of  the  crystals 
in  a  test  tube  with  a  little  hydrochloric  acid  and  warm  gently. 

Identify  the  gas  liberated  (Hint:  Odor?).  Explain.  What  is  the 
green  substance  formed  which  remains  in  solution? 

2CrO3  +  i2HCl3->  2CrCl3  +  30,  +  6H2O. 

Pour  a  drop  or  two  of  alcohol  upon  some  of  the  crystals.  The  action 
(oxidation)  is  so  violent  that  the  alcohol  frequently  takes  fire.  Try  the 
solubility  of  some  of  the  crystals  in  water  ( ?).  What  is  formed  ?  Equa- 
tion? 

The  oxidizing  power  of  chromates  and  dichromates  in  the  presence  of 
an  acid  is  due  to  the  presence  of  what  compound? 

Experiment  VII. — Formation  of  Perchromic  Acid. 

To  5  cm.3  of  a  solution  of  potassium  chromate  add  sulphuric  acid  until 
the  yellow  color  of  the  solution  changes  to  orange  (avoid  large  excess  of 
acid),  then  add  hydrogen  peroxide.  If  the  procedure  has  been  successful, 
blue  color  will  be  imparted  to  the  solution  by  the  unstable  perchromic  acid, 
H2Cr208. 

*  If  the  solutions  are  dilute  and  there  is  not  present  a  body  capable  of  being  oxid- 
ized the  decomposition  is  limited  to  the  liberation  of  chromic  acid,  H2O.Cr  O3. 


ELEMENTS    OF    GROUP    VI.  347 

Note. — The  perchromic  acid  is  supposed  to  be  a  compound  of  CrO3 
and  H2O2. 

K2CrO4  +  H2SO4->  (H2O.CrO3)  +  K2SO4. 

^  (H2O.CrO3)  ->  K2Cr2O7.+  H2O. 
2K2Cr~64  +  lp"O4-»  K2Cr2O7  +  K2SO4  +  H2O, 

K2Cr207  +  H2S04  —  K2S04  +  2CrO3  +  H2O 
2Cr03  +  H202  —  H2Cr2Os. 

The  above  reaction  is  used  as  a  delicate  test  for  either  hydrogen  perox- 
ide or  a  chromate. 

Experminent  VIII. — Chromium  as  a  Base-forming  Element.  Re- 
duction of  a  Chromate  to  a  Chromic  Salt.  Chrome  Alum. 

To  a  cold  saturated  solution  of  potassium  dichromate  acidified  with 
sulphuric  acid,  add  any  one  of  the  following  reducing  agents  until  the 
reddish-yellow  color  has  changed  to  a  dark  violet:  SO2  (sulphurous 
acid),  H2C2O4.2H2O  (oxalic  acid),  C2H5OH  (alcohol). 

The  solution  now  contains  chromium  sulphate,  Cr2(SO4)3.  The  color 
is  due  to  the  presence  of  what  ion  ?  Set  the  solution  aside  and  allow  it  to 
evaporate  spontaneously.  Observe  the  color  and  form  of  the  crystals. 
Compare  them  with  crystals  of  chrome  alum,  K2SO4,Cr2(SO4)3.24H2O. 
Your  conclusions  as  to  the  identity  of  the  crystals  ?  The  changes  may  be 
represented  thus: 

2Cr03  +  3S02  ->   Cr2(S04)3, 

K2Cr207  +H2O.S02+  H2SO4—  Cr2(SO4)3  +  K2SO4  +  4H2O. 
K2Cr207  +  4H2S04->  Cr2(S04)3  +  K2S04  +  (3O)  +  4H2O, 
3C2H5OH  +  (30)  ->  3C2H40  +  3H20. 

(Aldehyde) 
K2SO4  +  Cr2(SO4)3  +  24H2O  — (K2SO4,Cr2(SO4)3.24H2O) 

Chrome  alum. 

Is  chrome  alum  a  double  or  complex  salt  ?  Give  reasons  for  your  an- 
swer. Does  chromium  function  in  this  salt  as  an  acid  or  base-forming 
element  ?  Potassium  chrome  alum  is  the  analogue  of  potassium  aluminum 
sulphate  ("alum").  How  can  a  chromate  be  changed  to  a  dichro- 
mate ?  A  dichromate  to  a  chromate  ?  A  chromate  to  a  chromic  salt  ? 

Experiment  IX. — Hydrolytic  Decomposition  of  Chromium  Sulphate. 

Place  5  cm.3  of  a  solution  of  chromium  sulphate  (chrome  alum  will  do) 
in  a  test  tube  and  warm  gently.  Note  the  transition  from  the  violet 
chromic  sulphate  to  the  green  chromic  sulphate.  This  transition  has 
been  attributed  to  hydrolytic  decomposition: 

2Cr2(S04)3  +  H20^±  (Cr40(S04)4)S04  +  H2SO4. 
Experiment  X. — Formation  of  Chromyl  Chloride.   A  Test  for  a  Chloride. 

In  a  test  tube  fitted  with  a  delivery  tube  warm  a  mixture  of  a  chloride 
and  potassium  dichromate  with  strong  sulphuric  acid.  Pass  the  red- 


348  EXPERIMENTAL  CHEMISTRY. 

brown  vapor  (chromyl  chloride,  CrO2Cl2)  which  is  disengaged  into  a 
second  test  tube  containing  a  solution  of  an  alkaline  hydroxide.  The 
formation  of  a  chromate  is  indicated  by  the  yellow  color  which  the  solution 
assumes,  and  may  be  confirmed  by  acidifying  and  adding  a  lead-nitrate 
solution.  Presence  of  the  chromate  is  proof  of  the  presence  of  a  chloride 
in  the  original  mixture.  As  no  corresponding  bromine  and  iodine  com- 
pounds are  known,  it  is  obvious  that  by  means  of  this  test  it  is  possible 
to  detect  a  chloride  in  the  presence  of  either  a  bromide  or  iodide. 


K2Cr207  +  4KC1  +  3H2S04—  2Cr02Cl2  +  3K2SO4  +  3H,O. 

CrO2Cl2  +  4NH4OH  —  (NH4)2CrO4  +  2NH4C1  +  2H2O. 

Experiment  XI.  —  Formation  of  Chromic  Hydroxide. 

To  10  cm.3  of  a  solution  of  chrome  alum  add  ammonium  hydroxide 
in  slight  excess.  Result  ?  Equation  ?  Wash  the  precipitate  of  chromic 
hydroxide,  Cr(OH)3,  with  hot  water;  dry  and  reserve  it  for  the  following 
experiment. 

Experiment  XII.  —  Oxidation  of  a  Chromic  Compound  to  a  Chromate. 

(a)  Mix  a  portion  of  the  chromic  hydroxide  prepared  in  the  foregoing 
experiment  with  equal   portions  of   potassium  nitrate  and  sodium  car- 
bonate and  fuse  the  mixture  in  an  iron  crucible  or  on  a  piece  of  platinum 
foil.     Dissolve  the  fused  mass  in  water  and  filter.     What  is  the  color  of 
the  filtrate  ?     Test  it  for  potassium  chromate  by  acidifying  with  acetic 
acid  and  adding  a  solution  of  lead  nitrate.     Result? 

(b)  To  5  cm.3  of  a  solution  of  a  chromium  salt  (chrome  alum)  add 
sufficient  sodium  hydroxide  to   make   solution  alkaline.     How  add  a 
large  volume  of  bromine  water  and  heat  gently.     What  is  the  nature  of 
the  change?     Lead  peroxide  and  other  oxidizing  agents  may  be  sub- 
stituted for  the  bromine  water. 

Experiment  XIII.  —  Analytical  Reactions. 

(a)  Make  a  borax  bead,  touch  it  with  a  small  quantity  of  any  chromium 
compound  and  heat  in  both  the  oxidizing  and  reducing  flames.     The 
grass-green  color  is  imparted  to  the  bead  by  all  of  the  chromium  com- 
pounds. 

(b)  Reactions  of  chromates. 

1.  Recall  the  action  of  sulphuric  acid  on  a   solution  of  potassium 
chromate.     Equation? 

2.  Boil    a   potassium   dichromate  solution  which  has  been  acidified 
with  dilute  acid  with  i  cm.3  of  alcohol.     Account  for  results.     Equation  ? 

3.  Pass  hydrogen  sulphide  into  an  acid  solution  of  potassium  chro- 
mate.    Results  ?     Does  the  hydrogen  sulphide  act  as  an  oxidizing  or  re- 
ducing agent?     Equation? 

4.  Recall  the  action  of  hydrogen  peroxide  on  an  acid  solution  of  a 
chromate.     Equation  ? 

5.  Add  a  solution  of  barium  chloride  to  a  potassium  chromate  solu- 
tion (?).     Equation? 


ELEMENTS    OF    GROUP    VI.  349 

6.  Repeat    5,    using    a    potassium    dichromate    solution.     Compare 
results.     Equation  ? 

7.  Repeat  5  and  6,  using  a  solution  of  lead  nitrate  or  lead  acetate. 
Do  dichromates  precipitate  dichromates  or  chromates  when  added  to 
a  solution  of  a  salt  whose  chromate  is  insoluble? 

(c)  Reactions  of  chromic  salts. 

1.  Recall  the  effect  of  boiling  a  solution  of  a  chromic  salt  (see  Exp. 
IX).     Equation? 

2.  To    a    solution  of    potassium  chromium  sulphate   (chrome  alum) 
add  slowly  a  little  of  a  solution  of  sodium  hydroxide  ( ?) ;  then  in  excess  ( ?) 
Equations?     Does  aluminum   hydroxide   behave   in   a   manner   similar 
to  the  chromium  hydroxide  ?     Equation  ?     Is  the  chromium  hydroxide 
reprecipitated  when  the  solution  is  boiled? 

3.  Treat  a  solution  of  chrome  alum  with  sodium  carbonate  solution. 
Account  for  the  precipitation  of  chromium  hydroxide.     Ionic  equations  ? 

4.  Add  ammonium  sulphide  to  another   portion    of    the  solution  of 
chrome  alum.     Filter,  and  wash  the  precipitate  until  free  from  odor. 
Prove  that  it  is  not  a  sulphide.     What  other  metal  fails  to  form  a  stable 
sulphide  in  the  presence  of  water  ?     Equation  ? 

(d)  Ions.     What  ions  are  yielded  by  chromates  and  chromic  salts? 

MOLYBDENUM,    MO. 

At.  Wt.  96.0         Sp.  Gr.  8.6. 

Molybdenum  is  found  chiefly  in  the  two  uncommon  minerals,  molyb- 
denite, MoS2,  and  wuljenite,  PbMoO4.  It  is  a  hard,  silver-wrhite  metal 
which  is  less  fusible  than  platinum.  At  ordinary  temperatures,  it  re- 
mains unchanged  in  the  air,  but  on  heating  it  oxidizes  to  molybdenum 
trioxide,  MoO3.  This  oxide  interacts  with  the  alkalies  forming  molyb- 
dates  of  the  type  of  sodium  molybdate,  Na2MoO4.ioH2O.  Nitric  acid 
precipitates  a  hydroxide  MoO(OH)4  from  solutions  of  these  molybdates. 
This  hydroxide  gives  molybdic  acid,  H2MoO4,  when  heated  sufficiently 
to  dry  it.  When  sodium  phosphate  is  added  to  a  solution  of  ammonium 
molybdate  in  nitric  acid  the  yellow  pulverulent  ammonium  pliosphomolyb- 
date,  (NH4)3PO4.iiMoO3.6H2O,  is  precipitated.  This  is 'one  of  the 
more  familiar  of  the  complex  phosphomolybdates.  The  formulae  of  some 
of  the  more  important  compounds  of  molybdenum  are  as  follows:  MoO, 
Mo2O8,  MoO2,  MoO3,  MoCl2,  MoCl3,  MoCl4,  MoCl5,  MoO2Cl2,  MoOCl4, 
MoS2,  MoS3,  Na2MoO4,  K2Mo3O13.ioH2O,  Na2Mo4O13.6H2O. 

It  is  evident  that  molybdenum  possesses  valences  ranging  ordinarily 
from  II  to  VI. 

TUNGSTEN,    W. 

At.  Wt.  184.0         Sp.  Gr.  19.1. 

Tungsten  occurs  in  the  minerals  wolframite,  3FeWO4.3MnWO4, 
scheelite,  CaWO4,  and  hubnerite,  MnWO4.  The  metal  possesses  prop- 


350  EXPERIMENTAL  CHEMISTRY. 

erties  similar  to  those  of  molybdenum.  If  forms  a  class  of  compounds, 
many  of  which  are  analogues  of  the  compounds  of  the  previously  men- 
tioned element.  The  salts  corresponding  to  tungstic  acid,  H2WO4,  are 
known  as  tungstates.  The  following  list  contains  the  formulae  of  some 
of  the  better  known  compounds:  WO2,  WO3,  WC12,  WC14,  WC15,  WCle, 


The  metal  is  used  in  the  manufacture  of  tungsten  steel,  a  very  hard 
variety  of  which  contains  about  5  per  cent,  of  tungsten. 

URANIUM,    U. 

At.  Wt.  238.5         Sp.  Gr.  18.7. 

Uranium  is  found  in  the  ore  pitchblende,  which  contains  the  mineral 
uraninite,  U3O8,  and  several  other  rare  minerals.  It  is  a  heavy,  silvery- 
white  metal  which  decomposes  water  at  ordinary  temperatures  and  burns 
in  the  air  at  175°  when  in  the  powdered  form.  Some  of  the  more  im- 
portant compounds  are,  uranous  oxide,  UO2,  uranic  anhydride,  UO3, 
uranous  chloride,  UC14,  uranic  or  uranyl  sulphate,  UO2SO4.6H2O, 
sodium  diuranate,  Na^C^Oy.  Many  of  the  uranyl  compounds  are  yel- 
low in  color,  with  green  fluorescence.  Sodium  diuranate  is  used  in  mak- 
ing uranium  glass  which  shows  a  yellowish-green  fluorescence. 

The  compounds  of  uranium  are  of  particular  interest  because  of  being 
the  object  of  much  investigation  in  connection  with  the  phenomena  of 
radio-activity.  In  1898,  Becquerel  noticed  that  all  compounds  of  uran- 
ium gave  out  radiation  capable  of  affecting  a  photographic  plate  covered 
with  dark  light-proof  paper.  His  observations  eventually  led  to  the 
discovery,  by  Mme.  Curie  of  the  element  radium.  (The  student  is  urged 
to  consult  some  text-book  on  the  subject  of  radio-activity,  for  example, 
Rutherford,  —  Radio-activity.) 

PROBLEMS. 

1.  Calculate    the    percentage    composition    of    (a)    lead    chromate, 
PbCrO4,  (b)  chromite,  FeCr2O4. 

2.  If  10  grams  of  chromous  chloride,  CrCl2,  yield  5.75  grams  of  chlor- 
ine, what  is  the  atomic  weight  of  chromium? 

3.  10  grams  of  silver  chloride  are  formed  by  the  interaction  of  silver  ni- 
trate and  3.6865  grams  of  chromic  chloride;  what  is  the  atomic  weight 
of  chromium? 


3AgN03  +  CrCl3—  Cr(N03)3 


CHAPTER  XXXII. 


ELEMENTS  OF  GROUP  VII. 

Family  M. 
MANGANESE,  Mn. 

In  the  periodic  table,  according  to  the  present  classification,  man- 
ganese stands  alone  on  the  left  side  of  the  eighth  column.  The  elements 
of  the  halogen  group  occupy  the  right  side.  It  is  both  an  acid-forming 
and  base-forming  element,  and  probably  forms  as  large  a  variety  of 
compounds  as  any  element  known.  This  latter  property  is  due  to  the 
many  degrees  of  valence  which  manganese  can  manifest. 

Although  manganese  occurs  native  in  small  amounts  associated  with 
iron  in  meteorites,  it  is  found  principally  in  combination  in  minerals, 
some  of  which  are  rather  widely  distributed.  The  chief  source  of  the 
metal  is  pyrolusite,  MnO2.  Other  minerals  containing  it  are:  baunite, 
Mn2O3,  hausmannite,  Mn3O4,  the  hydrated  form,  manganite,  MnO(OH), 
manganese  spar,  MnCO3,  and  manganese  blende,  MnS. 

Manganese  may  be  obtained  by  heating  the  oxides  with  carbon  in 
an  electric  furnace  or  by  electrolysis  of  the  fused  chloride.  It  is  prepared 
more  conveniently,  however,  by  the  "  Goldschmidt  Process";  i.  e.,  by 
mixing  the  oxide  with  finely  divided  aluminum  and  igniting  the  mixture. 
The  aluminum  takes  the  oxygen  and  sets  free  the  manganese. 

Manganese  is  a  hard,  grayish-white,  brittle  metal  of  brilliant  luster. 
It  is  slightly  magnetic  and  fuses  at  1900°.  It  is  permanent  in  dry  air, 
but  readily  oxidizes  superficially  on  exposure  to  moist  air,  and  slowly 
decomposes  boiling  water  with  the  evolution  of  hydrogen  when  the 
metal  is  in  the  finely  divided  condition.  It  dissolves  readily  in  dilute 
acids  with  the  formation  of  manganous  salts.  The  metal  is  used  as  a  com- 
ponent of  the  alloys,  ferro-manganese  (20  to  75  per  cent,  manganese), 
and  spiegeleisen  (iron-manganese  carbide),  which  are  of  use  in  the  metal- 
lurgy of  Bessemer  steel.  " Manganin"  is  an  alloy  of  copper,  manganese 
and  nickel,  containing  from  eight  to  ten  per  cent,  of  manganese  and  three 
to  four  per  cent,  of  nickel.  In  the  form  of  wire  it  is  much  used  in  the 
construction  of  resistance  coils  because  of  its  very  low  resistance  tem- 
perature coefficient. 

Manganese  forms  five  rather  well  defined  sets  of  compounds  which 
correspond  to  its  series  of  oxides.  The  composition  of  these  compounds 
is  as  follows: 


Manganous, 
MnO 
Mn(OH)2 
MnSO4 
MnCl2 
etc. 


Manganic, 

Mn2O3 

Mn(OH)3 

Mn2(S04)3 

(MnCla) 


Manganites, 
MnO2 
H2MnO3 
CaMnO3 

etc. 


Manganates, 
MnO3 
H2MnO4 
K2MnO4 
etc. 


Permanganates. 
Mn2O7 
HMnO4 
KMnO4 
etc. 


352  EXPERIMENTAL  CHEMISTRY. 

The  manganous  salts  are  pale  pink  in  color  and  are  but  slightly  hydro- 
lyzed.  The  oxide  is  green  powder  and  a  strong  base.  The  manganic 
salts  are  violet  in  color  and  are  completely  hydrolyzed.  Manganese 
sesquioxide  is  a  weak  base.  The  manganites  are  usually  of  dark  color 
and  are  strongly  hydrolyzed.  Manganese  dioxide  is  a  black  solid  and 
behaves  as  an  indifferent  oxide.  The  manganates  are  green  in  color 
and  are  very  easily  hydrolyzed,  the  free  acid  decomposing  and  yielding  a 
higher  acid  (HMnO4)  and  a  lower  oxide  (MnO2). 

Manganese  trioxide  is  an  amorphous  red  solid  soluble  in  wrater.  It  is 
viewed  as  the  anhydride  of  manganic  acid,  H2MnO4.  The  permanganates 
are  purplish-red  in  color  and  are  not  hydrolyzed  by  water.  The  sept- 
oxide  of  manganese,  Mn2O7,  is  an  oily,  dark  liquid  which  is  presumably 
the  anhydride  of  permanganic  acid,  HMnO4. 

Manganese  has  a  valence  of  II,  III,  IV,  VI  and  VII.  If  the  existence 
of  manganese  tetroxide,  MnO4,  is  admitted,  then  the  metal  has  a  maxi- 
mum valence  of  VIII. 

It  should  be  noted  that  as  the  valence  of  manganese  increases,  its 
basic  nature  diminishes,  and  that  it  loses  all  of  its  basic  nature  and 
manifests  strong  acid  properties  wrhen  in  the  condition  of  higher  valence. 

MANGANESE,  MR. 

At.  Wt.  55.0     Sp.  Gr.  7.2-8.0. 

Experiment  I. — Note  the  most  obvious  physical  properties  of  man- 
ganese. 

Does  it  oxidize  in  the  air?     What  are  some  of   the  uses  of  manganese? 
Experiment  II. — Compounds  of  Manganese. 

Follow  directions  given  in  Experiment  II,  "  Chromium,"  substituting 
the  word  permanganate  for  the  word  dichromate. 

Experiment  III. — Preparation  of  a  Manganate.  Oxidation  of  a 
Manganate  to  a  Permanganate. 

(a)  Grind  in  a  mortar  5  grams  of  potassium  hydroxide  and  2.5  grams 
of  potassium  chlorate;  transfer  the  mixture  to  an  iron  crucible  (T.  O.) 
or  a  porcelain  crucible,  and  heat  until  the  mixture  fuses,  then  add  gradu- 
ally 5  grams  of  powdered  manganese  dioxide  while  stirring  with  an  iron 
rod  or  the  reverse  end  of  a  file.     Maintain  the  crucible  at  a  red  heat  for 
15  to  20  minutes.     Dissolve  the  green  mass  (potassium  manganate)  in  a 
little  cold  water,  then  decant  the  clear  liquid  away  from  the  residue. 
What  is  the  color  of  the  solution?     This  color  is  due  to  the  presence 
of  what  ion? 

2  MnO2  +  6KOH  +  KC1O3  ->  3K2MnO4  +  KC1  +  3H2O- 

(b)  Divide  the  green  solution  from  (a)  into  four  parts.     Dilute  one 
portion  with  an  equal  volume  of  distilled  water,  and  set  aside  for  several 


ELEMENTS    OF    GROUP    VII.  353 

days.  It  will  become  purple  in  color  owing  to  the  oxidation  of  the 
manganate  to  the  permanganate.  Pass  carbon  dioxide  (generator) 
through  one  portion  of  the  manganate  solution.  Result? 

3K2MnO4  +  2CO2->  2KMnO4  +  MnO2  +  2K2CO3. 

What  is  the  color  of  the  permanganate  ion?     Interpret  the  following 
equations: 

K2MnO4  +  2H2O  ->  2KOH  +  (H2MnO4),  (i) 

or  2K'  +  MnO"4  +  2H'  +  2OH'->  2K'  +  2OH'  +  (H2MnO4), 

H2MnO4  =  H2O,MnO3, 

3(H2O,MnO3)  —  H2O,Mn2O7  +  MnO2  +  2H2O,   (2) 
or  H2O,Mn2O7  =  2HMnO4 

3H2MnO4—  2HMnO4  +  MnO2  +  2H2O         (3) 
2KOH  +  2HMnO4—  2KMnO4  +  2H2O 

The  partial  equations  may  be  summarized: 

3K2MnO4  +  2H2O  —  2KMnO4  +  4KOH  +  MnO2 

The  ionic  equation: 

6K'  +  3MnO"4  +  2H'  +  2OH'->  6K'  +  2MnO/4  +  Mn  O2  +  4OH'. 

What  is  the  difference  between  the  manganate-ion  and  the  perman- 
ganate-ion ? 

Experiment  IV. — Oxidizing  Power  of  Manganates  and  Permanganates. 

Note.— Potassium   permanganate   as   well   as   potassium   manganate 
are  powerful  oxidizing  agents. 

(a)  Manganates.     To  a    warm  dilute   solution  of    oxalic  acid    add 
an  alkaline  solution  of  potassium  manganate  drop  by  drop.     The  oxalic 
acid  is  oxidized  to  carbon  dioxide  (prove  that  it  is  evolved)  and  water. 
The  manganese  is  reduced  to  the  bivalent  condition  with  the  formation 
of  manganous  sulphate,  MnSO4. 

(b)  Permanganates. 

1.  Repeat  (a)  with  a  solution  of  potassium  permanganate  to  which 
has   been   added   an   equal   volume  of   dilute   sulphuric  acid.     Result? 
Equation  ? 

2.  To  5  cm.3  of  sulphurous  acid  add  an  acidified  solution  of  potassium 
permanganate  drop  by  drop.     Results  ?     Equation  ? 

3.  Add   sulphuric   acid   to    a   solution    of   ferrous    sulphate,   FeSO4, 
then  add  potassium  permanganate  solution  drop  by  drop.     The  ferrous 
sulphate  is  oxidized  to  ferric  sulphate,  Fe2(SO4)3,  and  the  permanganate 
is  reduced.     Explain  the  color  changes.     Equation? 

4.  Add  a  solution  of   potassium   hydroxide   to  a  potassium  perman- 
ganate solution.     Result?     Equation? 

5.  Recall  or  again  try  the  action  of  hydrogen  peroxide"  on  an  acidified 
solution  of  potassium  permanganate. 


354  EXPERIMENTAL  CHEMISTRY. 

Interpret  the  following  equations: 

2KMnO4  +  3H2SO4-*K2SO4  +  2MnSO4  +  3H2O  +  (50). 

5C2H204  +  (50)  -r-  ioC02  +  5H20. 
2FeS04  +  H2S04  +  (O)  —  Fe2(SO4)3  +  H2O. 

H2SO3  +  (O)  —  H2SO4 

4KMnO4  +  4KOH  —  4K2MnO4  +_2_HaO  +  O3 
2KMnO4  +  (heat)  —  K2MnO4  +  O2. 

Experiment  V. — Reactions  of  Manganous  Salts. 

Note. — Use  a  solution  of  manganous  chloride,  MnCl2,  in  performing 
the  following  experiments. 

(a)  What  color  is  imparted  to  the  borax  bead  by  manganese  com- 
pounds when  heated  in  the  oxidizing  flame  ?     In  the  reducing  flame  ? 

(b)  Mix   a   little   manganous   chloride   with    sodium   carbonate   and 
potassium  nitrate,  and  fuse  on  a  platinum  foil.     What  is  the  color  of 
the  fused  mass?     Identify  the  substance  (?). 

(c)  (i)  To  a  dilute  solution  of  manganous  chloride  add  ammonium 
hydroxide.     Result?     Equation? 

(2)  Repeat  (i)  using  a  solution  of  manganous  chloride  to  which  has 
been  added  ammonium  chloride.  Result  ? 

(d)  Repeat  (c)  (i),  using  sodium  hydroxide  solution  in  excess.     Equa- 
tion?    Treat  the  manganous  hydroxide  with  bromine  water. 

(e)  Repeat  (c),  using  a  solution  of  sodium  carbonate  or  ammonium  car- 
bonate.    Result?    Prove  that  the  precipitate  is  a  carbonate  (?).     Equa- 
tions ? 

(/)  Add  ammonium  sulphide  to  a  solution  of  manganous  chloride. 
Prove  that  the  precipitate  is  a  sulphide  (?).  Equations? 

(g)  Pass  hydrogen  sulphide  (generator  in  hood)  through  a  solution 
of  manganous  chloride  to  which  has  been  added  i  cm.3  of  acetic  acid. 
Account  for  the  negative  results. 

(h)  What  is  the  color  of  the  manganous  ion  ?  Of  the  manganate  ion  ? 
Of  the  permanganate  ion? 

PROBLEMS. 

i.  A  liter  of  a  solution  of  potassium  permanganate  contains  10  grams 
of  the  salt.  How  many  grams  of  ferrous  sulphate,  FeSO4,  can  this 
solution  oxidize  to  ferric  sulphate,  Fe2(SO4)3,  in  the  presence  of  the 
proper  amount  of  sulphuric  acid? 


CHAPTER  XXXIII. 
TRANSITION  ELEMENTS. 

FIRST    LONG    PERIOD    (IRON    ELEMENTS). 

Iron,  Fe.  55.9 
Cobalt,  Co.  59.0 
Nickel,  Ni.  58.7 

The  iron  elements  are  among  the  most  important  elements  tech- 
nically, and  are  most  interesting  from  the  chemical  stand-point. 

These  three  elements,  iron,  cobalt  and  nickel,  stand  in  a  different 
relation  to  one  another  than  the  members  of  the  other  eight  groups;  i.e., 
they  are  not  the  corresponding  members  of  successive  periods  as  are 
the  elements  of  the  families.  They  belong  to  the  same  period  and  form  a 
transition  group  between  the  first  and  second  series  of  the  first  long 
period.  (See  Periodic  Classification.)  These  three  elements  are 
closely  related;  in  nature  they  are  usually  associated;  the  free  metals 
are  magnetic  and  possess  many  physical  properties  in  common.  They 
exhibit,  however,  a  gradual  transition  in  their  chemical  properties. 
Thus,  iron  forms  ferrates,  M'2FeO4,  and  two  basic  oxides,  ferrous  oxide, 
FeO,  and  ferric  oxide,  Fe2O3,  each  of  which  yields  a  series  of  stable  salts 
of  the  type,  FeCl2  (ferrous  chloride)  and  FeCl3  (ferric  chloride),  respec- 
tively. Cobalt  forms  cobaltous  and  cobaltic  salts,  like  CoCl2  and  Co2(SO4)3. 
Many  of  the  cobalt^  salts  (except  the  double  salts)  are  unstable.  Nickel 
forms  only  one  series  of  salts,  namely,  mckelous  salts,  like  NiCl2. 

It  is  obvious  that  these  metals  are  related,  on  the  one  hand,  through  iron, 
to  chromium  and  manganese  (recall  such  compounds  as  chromates, 
manganate  and  ferrates),  and  on  the  other  hand,  through  nickel,  to  copper 
and  zinc,  both  of  which  are  bivalent  elements  and  follow  in  the  period. 

IRON,  Fe. 
At.  Wt.  55.9         Sp.  Gr.  7.78. 

Iron  is  one  of  the  most  abundant  and  widely  distributed  elements, 
yet  it  is  found  native  in  only  small  amounts  in  meteorites  and  certain 
volcanic  ejects.  Most  of  the  rocks  contain  compounds  of  iron,  and 
many  of  the  red  and  yellow  soils  owe  their  color  to  the  presence  of  iron 
compounds.  Mineral  waters  holding  iron  compounds  in  solution  are 
known  as  "chalybeate  waters."  Minute  .quantities  of  compounds  of 
iron  are  also  found  in  chlorophyl  and  in  the  haemoglobin  of  the  blood.* 

*  Ammonium  sulphide  interacting  with  the  iron  compounds  present  in  the  tissues 
blackens  the  skin 

355 


356  EXPERIMENTAL  CHEMISTRY. 

The  chief  ores  of  iron  are  red  hematite  and  specular  iron  ore,  Fe2O3  (ores 
found  in  the  Lake  Superior  region,  particularly  in  northern  Michigan, 
in  Alabama,  in  Missouri  and  other  regions  of  the  United  States) ;  limonite 
or  brown  hematite,  2Fe2O3,3H2O  (ores  found  chiefly  in  Alabama  and 
several  other  Southern  States);  magnetite  or  magnetic  iron  ore  (loadstone), 
Fe3O4  (ores  found  in  Pennsylvania,  New  York,  Michigan  and  New 
Jersey);  jranklinite,  similar  to  magnetite;  and  siderite  or  spathic  iron  ore, 
FeCO3.  Iron  is  also  found  in  combination  wTith  sulphur  as  iron  pyrites 
(Fool's  gold),FeS.  This  pyrite  is  used  in  the  manufacture  of  sulphuric 
acid.  Chalcopyrite,  a  sulphide  of  iron  and  copper,  contains  small  quan- 
tities of  cobalt  and  nickel. 

Iron  ores  are  usually  reduced  by  heating  them  with  carbon  and  a  flux 
in  a  blast- furnace.  The  nature  of  the  flux  depends  on  the  composition 
of  the  iron  ores;  if  the  ores  contain  silica  and  clay,  a  flux  of  basic  nature, 
like  limestone,  is  used;  and,  conversely,  ores  mixed  with  lime  or  magnesia 
are  heated  with  an  acid  flux,  such  as  sand  or  clay-slate,  in  order  that  a 
fusible  slag  may  be  formed.  An  impure  iron  (pig  iron)  is  obtained  by 
this  process. 

(Discussions  of  the  composition,  properties  and  uses  of  cast  iron, 
wrought  iron,  Spiegel  iron  and  steel  will  be  found  in  the  reference  texts. 
The  student  is  urged  to  become  familiar  with  the  principles  involved  in 
the  manufacture  of  steel  by  such  methods  as  the  "Bessemer  process," 
the  " Thomas-Gilchrist  process,"  and  the  "Siemens-Martin  process" 
or  "  Open-hearth  process.") 

Pure  iron  may  be  prepared  by  reducing  the  oxide  or  oxalate  in  a 
stream  of  hydrogen;  electrolytic  iron  may  be  deposited  from  solutions 
of  certain  salts  by  a  properly  regulated  electric  current;  or  the  pure  metal 
may  be  prepared  by  Goldschmidt's  process. 

It  is  a  white,  lustrous  metal,  ductile  and  more  malleable  than  wrought 
iron.  When  finely  divided  it  has  a  gray  color.  It  is  very  tenacious; 
and  it  is  to  this  property,  together  with  its  abundance  and  ease  with  which 
it  can  be  prepared,  that  makes  it  the  most  valuable,  industrially,  of  all 
metals.  Pure  iron  melts  at  1800°;  wrought  iron,  at  1600°;  and  cast  iron, 
noo°-i3oo°.  At  red  heat  it  becomes  soft  and  can  be  welded.  It  is 
attracted  to  a  magnet,  but  does  not  retain  its  magnetism.  Iron  is  not 
acted  upon  by  dry  air  at  ordinary  temperatures,  but  in  moist  air  it  be- 
comes coated  with  a  "rust"  which  is  probably  a  mixture  of  the  oxide 
and  the  hydroxide  of  iron,  2Fe2O3,  (FeOH)3.  It  is  not  definitely  under- 
stood just  how  the  product*  is  formed.  The  alkali  hydroxides  or  car- 
bonates prevent  rusting.  At  red  heat,  massive  iron  decomposes  water, 
while  finely  divided  iron  decomposes  water  at  100°.  Iron  dissolves  in 
dilute  hydrochloric  and  sulphuric  acids,  hydrogen  being  evolved.  Dilute 
nitric  acid  dissolves  iron  with  the  formation  of  ferrous  nitrate  and  am- 
monium nitrate;  concentrated  nitric  acid  yields  ferric  nitrate  and  oxides 
of  nitrogen.  When  steel  and  cast  iron,  which  contain  iron,  iron  carbide, 
Fe3C,  and  graphite  are  treated  with  cold  dilute  acids  almost  pure  hydrogen 
*The  Corrosion  of  Iron,  A.  S.  Cushman.  Bui.  No.  30,  Dep't  Agr. 


TRANSITION    ELEMENTS.  357 

is  liberated,  as  the  carbide  and  the  graphite  are  not  attacked;  more  con- 
centrated acids,  however,  decompose  the  carbide,  with  the  assistance 
of  nascent  hydrogen,  the  freed  carbon  combines  with  hydrogen  forming 
hydrocarbons,  which  mix  with  the  escaping  gas,  and  give  to  it  a  very 
disagreeable  odor.  With  nitric  acid,  the  carbide  carbon  behaves  in  a 
slightly  different  manner.  In  this  case  the  carbide  passes  into  solution 
in  combination  with  hydrogen  and  oxygen.  This  compound  imparts  a 
distinct  color  to  the  solution.  Under  similar  conditions  the  depth  of  the 
color  varies  with  the  proportion  of  carbon  present.  The  graphite,  how- 
ever, is  set  free  and  separates  out  as  with  other  acids.  When  a  piece  of 
iron  is  immersed  in  very  strong  nitric  acid  (sp.  gr.  1.42)  for  an  instant 
and  then  removed,  it  is  not  attacked  by  ordinary  nitric  acid  and  is  in- 
capable of  displacing  hydrogen  and  other  elements  lying  below  it  in  the 
electromotive  series.  Iron  exhibiting  these  peculiar  properties  is  said 
to  be  in  the  passive  state.  Such  iron  loses  these  properties  and  becomes 
active  when  it  is  struck  a  sharp  blow  or  is  brought  into  contact  with 
ordinary  iron.  It  was  formerly  supposed  that  the  apparent  chemical 
inertness  of  the  metal  was  due  to  the  formation  of  a  protecting  layer 
of  oxide  over  its  surface,  but  it  has  been  shown  recently  to  the  satisfaction 
of  many  that  the  passivity  is  due  to  an  electrical  condition  of  the  metal.* 
Nearly  all  of  the  compounds  of  iron  are  "ous"  compounds  in  which 
the  metal  is  a  bivalent  positive  ion,  or  "ic"  compounds  in  which  iron 
is  a  trivalent  positive  ion.  The  oxides  and  hydroxides,  FeO  and  Fe(OH)2, 
Fe2O3  and  Fe(OH)3,  are  basic,  the  ferrous  hydroxide  being  the  stronger 
base.  The  ferric  salts  derived  from  Fe(OH)3  are  hydrolyzed  to  con- 
siderable extent.  The  ferrous  salts  are  easily  oxidized  by  the  air  to  the 
ferric  condition.  Only  a  few  ferrates,  K2FeO4,BaFeO4,  are  known. 
Such  salts  as  potassium  ferrocyanide,  K4.Fe(CN)6  and  potassium  ferri- 
cyanide,  K3.Fe(CN)6,  yield  complex  anions  which  contain  this  element. 


IRON,    FC. 

At.  Wt.  55.9         Sp.  Gr.  7.8. 
Experiment  I. — Preparation  of  Iron  by  Reduction  of  Iron  Oxide. 

(a)  Reduction  by  hydrogen.  Assemble  apparatus  (Fig.  48)  as  described 
in  Chapter  IX,  Experiment  VIII.  Observe  the  precautions  suggested. 
Substitute  iron  oxide,  Fe2O3,  for  the  copper  oxide  in  the  porcelain  boat. 
Are  the  particles  of  iron  oxide  attracted  by  the  magnet  ?  After  the  glass 
tube  has  been  maintained  at  a  red  heat  for  10  to  15  minutes,  remove 
the  heat  and  allow  the  boat  and  contents  to  cool  in  a  stream  of  hydrogen. 
What  was  formed  in  the  anterior  portion  of  the  tube?  Examine  the 
substance  in  the  boat.  What  is  its  color?  Are  particles  of  it  attracted 
by  the  magnet?  Expose  it  to  the  air  (?).  Why  did  you  cool  the  sub- 

*  The  Polarization  Capacity  of  Iron  and  its  Bearing  on  Passivity. — Jour.  Amer. 
Chem.  Soc.,  Nov.,  1906. 


EXPERIMENTAL  CHEMISTRY. 


stance   in   an   atmosphere  of  hydrogen?     Identify  the   substance   (?). 
Equation  ?    What  is  "  Venetian  red  "  ? 

(b)  Goldschmidt's  method  (reduction  with  aluminum).  Mix  equal 
volumes  of  powdered  iron  oxide  (hammer  scale)  and  aluminum  powder, 
and  ignite  (Caution!)  the  mixture  on  an  iron  plate  or  a  brick.  This 


FIG.  48.     (Smith  and  Keller.) 

latter  operation  is  performed  most  readily  by  inserting  a  piece  of  mag- 
nesium ribbon  in  the  mixture  and  setting  fire  to  the  projecting  part. 
Result  ?  Equation  ? 

Note. — "Thermit"  may  be  substituted  for  the  above  mixture. 

Experiment  II. — Properties  of  Iron. 

(a)  Examine  cast  iron,   wrought  iron   and   steel.     Note  their  most 
obvious  physical  properties.     Give  the  approximate  composition  of  each. 
Examine  a  piece  of  piano  wire;  it  is  about  99.7  per  cent.  pure. 

(b)  Recall  the  behavior  of  a  piece  of  iron  heated  to  redness  when 
plunged  into  a  jar  of  oxygen.     Equation?     Recall  or   try  the  action  of 
iron  when  treated  with  dilute  acids.     Equations  ? 

Does  iron  "  rust "  in  dry  air  ?     In  moist  air  ?     Equation  ? 

Experiment  III. — Compounds  of  Iron. 

(a)  Examine  the  compounds  of  iron  (end  shelf).     Give  the  color  and 
structural  formula  of  each  compound.     Record  the  valence  of  the  iron 
atom  in  each  compound,  and  state  whether  it  manifests  the  properties 
of  an  acid-forming  or  base-forming  element.     Tabulate  the  above  data. 

(b)  What  is  the  formula,   color  and  electrical  charge  of  the  ferrous 
ion,  the  ferric  ion  and  the  f erro-cyanide  ion  ? 

Hint. — Examine  solutions  of  the  salts  which  yield  these  ions  (end 
shelf). 

Experiment  IV. — Preparation  of  Ferrous  Sulphate. 

Place  about  20  grams  of  iron  (free  from  rust)  in  the  form  of  filings, 
nails  or  wire,  in  an  Erlenmeyer  flask  in  which  there  should  be  inserted  a 
cork  fitted  with  a  glass  jet  to  allow  gas  to  escape.  Pour  on  the  iron  175 
cm.3  of  dilute  sulphuric  acid,  add  a  few  drops  of  concentrated  acid  and 


TRANSITION    ELEMENTS.  359 

warm  gently  if  the  action  appears  slow.  Note  the  odor  of  the  escaping 
gases  ( ?).  Account  for  the  odor.  Allow  the  action  to  continue  for  10  or 
15  minutes,  adding  sufficient  strong  acid  to  keep  up  a  brisk  action,  while 
other  experiments  are  proceeded  with.  When  nearly  (but  not  quite) 
all  of  the  iron  has  dissolved,  filter  into  a  casserole  containing  2  cm.3 
of  sulphuric  acid.  Note  the  color  of  the  solution  and  set  it  aside  for  a  day. 
Pour  off  the  mother  liquid  from  the  crystals,  and  wash  the  latter  with  cold 
water  by  decantation.  Dry  the  crystals  between  sheets  of  filter  paper. 
Note  their  color,  taste,  and  solubility  in  water  (?).  Put  a  few  of  them 
into  hard  test  tube  and  heat,  gently  at  first,  then  strongly  (?).  Compare 
with  the  corresponding  properties  of  ferrous  sulphate,  FeSO4.7H2O, 
crystals  (end  shelf).  Identify  the  prepared  crystals.  Equation?  Write 
the  structural  formula  for  ferrous  sulphate.  What  is  the  valence  of  iron 
in  this  compound? 

What  is  "green  vitriol"?  "oil  of  vitriol"?  "copperas"?  "white 
viriol "  ?  iron  protosulphate  ?  writing  ink  (black)  ? 

Experiment  V. — Preparation  of  Ferrous  Ammonium  Sulphate  (Mohr's 
Salt). 

Weigh  5  grams  of  ferrous  sulphate  into  an  evaporating  dish  or  a  cas- 
serole ;  calculate  the  weight  of  and  weigh  out  an  equi-molecular  quantity 
of  ammonium  sulphate.  Dissolve  the  salts  separately  in  the  smallest 
volume  of  hot  water.  Add  a  few  drops  of  sulphuric  acid  to  the  solution 
of  ferrous  sulphate.  Mix  the  solutions.  Allow  the  mixture  to  cool 
slowly  and  evaporate  spontaneously  (?).  Describe  the  crystals  and 
compare  them  with  those  in  the  laboratory  (end  shelf).  Collect  the 
crystals  in  a  funnel,  the  stem  of  which  is  closed  with  a  loose  plug  of 
glass-wool  (or  better,  provided  with  a  platinum  filter  cone),  allow  the 
mother  liquid  to  drain  off,  then  wash  the  crystals  with  a  small  quantity 
of  cold  water  and  dry  with  filter  paper. 

Is  "Mohr's  Salt,"  (NH4)2SO4,FeSO4,6H2O,  a  complex  or  compound 
salt  ?  Reason  for  your  answer  ?  Equation  ? 

Experiment  VI. — Preparation  of  Ferric  Ammonium  Sulphate,  Iron- 
Ammonium  Alum. 

Directions  are  the  same  as  those  given  in  Experiment  VI,  except  that 
ferric  sulphate,  Fe2(SO4)3,  is  used  instead  of  ferrous  sulphate.  Equation  ? 

To  which  class  of  salts,  complex  or  compound,  does  iron-ammonium 
alum,  (NH4)2SO4,Fe2(SO4)3,24H2O,  belong? 

Experiment  VII. — Hydrolysis  of  Ferric  Salts. 

Dissolve  equal  weights  (about  0.5  gram)  of  ferrous  sulphate  (use 
ferrous  ammonium  sulphate)  and  ferric  sulphate,  Fe2(SO4)3,  in  equal 
volumes  of  water  in  separate  test  tubes,  then  warm  slightly.  Observe 
the  color  of  each  by  looking  down  through  the  solution  at  a  piece  of  white 
paper  (?).  Test  each  solution  with  litmus  paper  (?).  Add  2  or  3  cm.3 
of  pure  concentrated  sulphuric  acid  to  each  solution  and  observe  the 


360  EXPERIMENTAL  CHEMISTRY. 

colors  again  (?).  What  is  the  color  of  the  ferrous  ion?  The  ferric 
ion  is  almost  colorless?*  Ferric  hydroxide,  Fe(OH)3,  is  a  reddishb-rown 
substance.  Can  you  account  for  the  change  in  color  of  the  solution  of 
the  ferric  salt?  Ionic  equations?  Which  manifests  the  stronger  basic 
properties — iron  in  the  "ous"  or  iron  in  the  "ic"  condition? 

Experiment  VIII.— Reactions  of  Ferrous  and  Ferric  Salts. 

(a)  Make  a  borax  bead  and  dissolve  in  it  a  small  quantity  of  any 
iron  compound  (preferably,  the  oxide);  treat  the  bead  successively  with 
the  oxidizing  (?)  and  reducing  flames  (?).  Use  a  recently  prepared 
solution  of  ferrous  sulphate  (ferrous  ammonium  sulphate)  and  a  dilute 
solution  of  ferric  chloride,  FeCl3,  for  the  following  reactions.  Treat  a 
portion  of  each,  separately,  with  the  following  reagents.  Compare  the 
two  results  obtained  with  each  reagent. 

(6)  Ammonium  hydroxide  (?).  Equations?  Allow  the  tubes  and 
their  contents  to  stand  exposed  to  the  air  for  15  or  20  minutes.  Result? 
Equations  ? 

(c)  Sodium  hydroxide  solution  (?).     Equation?     Do  the  precipitates 
dissolve  in  excess  of  the  precipitant? 

(d)  Sodium  carbonate  ( ?).     Prove  by  proper  tests  that  the  precipitates 
do  or  do  not  contain  the  carbonic  acid  radical.     Equations? 

(e)  A  few  drops  of  a  solution  of  potassium  ferricyanide,  K3Fe(CN)5  ( ?). 
Equation  ?     What  is  "  Turnbull's  blue  ?  " 

(/)  A  few  drops  of  a  potassium  ferrocyanide,  K4Fe(CN)3,  solution  (?). 
Equation  ?  What  is  "  Berlin  or  Prussian  blue  ?  " 

(g)  Several  drops  of  a  potassium  thiocyanate,  KCNS,  solution? 
Equation  ?  Inasmuch  as  the  ferric  ion  and  the  thiocyanate  ion  are  color- 
less, what  is  probably  the  source  of  the  color? 

(h)  What  reagents  would  you  use  to  test  for  the  presence  of  the  ferrous 
ion?  For  the  ferric  ion? 

Note. — Potassium  ferrocyanide,  potassium  ferricyanide  and  potas- 
sium thiocyanate  belong  to  a  class  of  substances  known  as  "  indicators"^ 
By  their  use  we  are  enabled  to  detect,  for  example,  the  presence  of  the 
iron  ion,  and  to  determine  also  its  state  of  oxidation;  i.e.,  whether  it  is 
in  the  "ous"  or  "ic"  condition.  The  use  of  potassium  thiocyanate 
constitutes  a  very  delicate  test  for  detecting  traces  of  the  ferric  ion. 

Experiment  IX. — Ferrocyanides  and  Ferricyanides. 

Using  diluted  solutions  of  potassium  ferrocyanide,  apply  tests  as  given 
in  Experiment  VIII,  (ft),  (c},  (d),  to  ascertain  whether  either  the  ferrous 

*  The  ferric  ion  is  almost  colorless.  The  yellowish-brown  color  of  solutions  of 
ferric  salts  is  due  to  the  piesence  of  ferric  hydroxide  (reddish-brown)  produced  by 
hydrolysis. 

t  See  Ostwald's  Scientific  Foundations  of  Analytical  Chemistry. 


TRANSITION    ELEMENTS.  361 

or  ferric  ions  are  present  or  not.  Result  ?  Is  potassium  ferrocyanide  a 
"complex"  or  a  "compound"  salt?  Why?  Indicate  by  an  equation 
how  it  probably  dissociates.  (See  note,  Chapter  XXV.  Experiment 
XVII.) 

Experiment  X. — Reactions   of   Iron   Salts   Involving  Oxidations   and 
Reductions. 

(a)  Action  of  hydrogen  sulphide  on  iron  salts. 

1.  Saturate  a  dilute  solution    of    a  ferrous    salt  (ferrous-ammonium 
sulphate)  with  hydrogen  sulphide  (?).     Add  ammonium  hydroxide  (?). 
Filter,  wash  thoroughly  the  precipitate  with  cold  water,  and  ascertain 
whether  it  is  a  sulphide  or  hydroxide.     Explain  in  terms  of  the  ion- 
product  constant  why  hydrogen  sulphide  does  not  not  precipitate  all  of 
the  iron  from  solutions  of  its  ferrous  salts.     Ferric  salts  are  reduced  by 
hydrogen  sulphide. 

2.  Pass  hydrogen  sulphide  into  a  dilute  solution  of  ferric  chloride  (?). 
Filter.     Test  the  precipitate  by  burning  it  on  a  piece  of  filter  paper  (note 
the  odor)  (?).     Test  the  filtrate  for  the  presence  of  ferrous  ions  (?). 

2FeCl3  +  H2S  —  2FeCl2  +  2HC1  +  S., 
2Fe'"  +  S"-»2Fe"  +_S. 

(b)  Action  of  ammonium  sulphide. 

1.  Ammonium    sulphide    precipitates    black   ferrous   sulphide,   FeS, 
from  solutions  of  ferrous  salts. 

2.  Ferric    salts    are   reduced    by   ammonium   sulphide.     The   latter 
produces  a  black  precipitate  of  ferrous  sulphide,  soluble  in  hydrochloric 
acid. 

2FeCl3  +  3(NH4)2S  -»  2 FeS  +  6NH4C1  +  S. 

(c)  Reduction  of  ferric  salts  to  ferrous  salts. 

1.  By  nascent  hydrogen.     Place  a  few  pieces  of  granulated  zinc  in  a 
small  flask  which  is  fitted  with  a  cork  provided  with  a  glass  jet  for  the 
escape  of  gases.     Pour  20  cm.3  of  ferric  chloride  into  the  flask  and  add  an 
equal  volume  of  water.    Now  add  sufficient  hydrochloric  acid  to  dissolve 
the  zinc  with  a  rapid  evolution  of  hydrogen.     The  action  is  hastened  by 
warming  gently.   Test  a  portion  of  the  solution  for  the  presence  of  ferrous 
ions  (?).     Is  the  ferric  ion  present  (test)  ?     Equations? 

2.  By  stannous  chloride.     To  3  cm.3  of  a  solution  of  ferric  chloride 
in  a  test  tube  add  i  cm.3  of  stannous  chloride.     Test  portions  of  the 
solution  as  in  (a),  i.     Equations? 

2FeCl3  +  SnCl2  — »  2FeCl2  +  SnCl4 
2Fe"'  +  Sn"  -*  2Fe"  +  Sn"" 

(d)  Oxidation  of  ferrous  salts  to  ferric  salts. 

i.  By  nitric  acid.     Into  5   cm.3  of   a   hot  dilute  solution  of  ferrous 


362  EXPERIMENTAL  CHEMISTRY. 

sulphate,  to  which  a  little  sulphuric  acid  has  been  added,  pour  5  cm.3 
of  concentrated  nitric  acid  drop  by  drop.  Test  separate  portions  of  the 
liquid  for  the  presence  of  ferrous  and  ferric  ions  (?).  Make  a  record 
of  all  tests  and  results.  Equations? 

2HN03  —  (30)  +  2NO  +  H20. 

6FeS04  +  3H2S04  +  2HNO3  —  3Fe2(SO4)3  +  4H2O  +  2NO. 
6FeCl2  +  3HC1  +  2HNO3  —  6FeCl3  +  4H2O  +  2NO. 

2.  By  potassium    bichromate   in    the  presence  of  an  acid.     Pour  5 
cm.3  of  dilute  sulphuric  acid  into  15  cm.3  of  ferrous  sulphate  solution  in 
a  small  beaker.     Now  add  by  means  of  a  burette  a  solution  of  potassium 
bichromate  drop  by  drop  until  a  drop  of  the  solution  of  iron  salt  trans- 
ferred to  white  porcelain  and  tested  with  potassium  ferricyanide  solution 
fails  to  give  a  blue  color.     This  test  shows  the  absence  of  what  ion? 
Test  drops  of  the  solution  with  a  solution  of  potassium  ferrocyanide  or  a 
solution  of  potassium  thiocyanate.     Result  ?     This  test  is  made  to  detect 
the  presence  of  which  ion  ?     Equations  ? 

K2Cr207  +  4H2S04->  K2S04  +  Cr2(SO4)3  +  4H2O  +  (3O). 
6FeS04  +  3H2S04  +  (3O)  —  3Fe2(SO4)3  +  3H2O. 

3.  By  potassium  permanganate  in  the  presence  of  an  acid.     Measure 
accurately  15  cm.3  of  a  recently  prepared  solution  of  ferrous  sulphate 
(ferrous-ammonium  sulphate)  into  a  small  clean  beaker,  and  add  5  cm.3 
of  dilute  sulphuric  acid.     Clamp  £  clean  burette  into  a  vertical  position 
and  fill  it  with  a  solution  of  potassium  permanganate  (end  shelf).     Allow 
this  latter  solution  to  drop  slowly  into  the  ferrous  salt.     The  pink  color 
of  the  permanganate  immediately  disappears  on   stirring  with  a  glass 
rod  (do  not  remove  the  rod  from  the  solution).     The  color  continues  to 
be  destroyed  until  all  of  the  ferrous  salt  is  completely  oxidized  to  the 
ferric   state,  when  a  drop  of  permanganate  added  in  excess  imparts  a 
faint  pink  color  to  the  liquid.     This  indicates  the  end  point]  i.e.,  that  the 
reaction  is  ended.     Record  the  number  of  cm.3  of  potassium  permanga- 
nate required  to  oxidize  the  ferrous  salt. 

By  applying  suitable  tests  to  portions  of  this  liquid,  prove  that  the 
oxidation  has  been  complete;  i.e.,  that  the  ferrous  ion  is  not  present 
and  that  ferric  ions  are  present.  Make  a  record  of  tests  employed  and 
results  secured. 

Repeat  the  experiment,  using  20  cm.3  of  ferrous  sulphate  solution. 
Equations  ? 

4.  Chlorine   water,   bromine   water,  and   potassium   chlorate  in   the 
presence  of   strong  hydrochloric  acid,  are  other  oxidizing  agents   fre- 
quently used  to  convert  ferrous  into  ferric  salts. 

5.  Add  5  cm.3  of  pure  concentrated  nitric  acid  to  50  cm.3  of  tap-water; 
evaporate  to  15-20  cm.3  and  test  for  the  presence  of  the  ferric  ion  (?). 


TRANSITION    ELEMENTS.  363 

COBALT,    CO. 

At.  Wt.  59.0         Sp.  Gr.  8.5. 

The  principal  ores  of  cobalt  are  smaltite,  CoAs2,  and  cobaltite,  CoAsS. 
The  pure  metal  may  be  obtained  by  reducing  the  oxide,  the  chloride 
or  the  oxalate  in  a  stream  of  hydrogen  or  by  Goldschmidt's  process. 
Cobalt  resembles  iron  in  many  respects.  It  is  a  lustrous  silver-white 
(pink-tinted),  hard  metal,  malleable,  tenacious,  and  when  heiated  is  very 
ductile.  It  melts  at  about  1500°.  Unlike  iron  and  nickel, t  retains  its 
magnetic  properties  even  at  red  heat.  The  metal  has  but  few  commer- 
cial applications,  its  use  being  confined  to  the  iron  and  steel  industry. 

Cobalt  in  the  massive  form  is  not  readily  acted  upon  by  the  air,  but 
the  finely  divided  metal,  especially  when  it  is  freshly  prepared  by  the 
reduction  of  the  oxide  in  hydrogen,  oxidizes  easily  and  may  take  fire 
spontaneously  in  the  air.  It  is  attacked  slowly  by  dilute  acids,  hydrogen 
being  liberated. 

Cobalt,  like  iron,  forms  two  kinds  of  ions — the  cobaltous  ion,  Co", 
and  the  cobaltic  ion,  Co'".  The  cobaltous  salts  are  but  slightly  hydro- 
lyzed,  but  the  cobaltic  salts  are  completely  decomposed  by  water.  The 
latter  are  rather  unstable  and  tend  to  break  down  into  the  cobaltous 
salts  with  a  liberation  of  one-third  of  the  acid  radical.  Most  of  the 
cobaltous  compounds  are  red  when  hydrated  and  blue  when  dehydrated. 
The  blue  color  of  the  salt  is  explained  by  some  chemists  as  being  due  to 
the  " repression  of  the  ionization  of  the  salt";  i.e.,  uto  the  driving  back 
of  the  ions  into  molecules,  which  are  blue."  Cobalt  shows  a  marked 
tendency  to  enter  into  combination  with  the  ions,  NO2,  CN  and  NH2, 
to  form  derivatives  which  yield  many  complex  ions  which  usually  give 
none  of  the  reactions  of  cobalt  ions.  The  cobalt  compounds  give  with 
ammonia  many  complex  compounds  which  present  some  very  complex 
relations.  One  of  the  most  interesting  series  of  these  complex  salts  is 
known  as  the  cobalt  amines. 

Some  of  the  more  important  compounds  are  represented  by  the  fol- 
lowing formulae: 

Cobaltous  compounds,  CoO,  Co(OH)2,  CoCl2,  6H2O,  CoSO4,6H2O, 
Co(NO3)2.6H2O,  CoCO3.6H2O,  CoS,  Co2(Fe(CN)6),  Co3(Fe(CN)6), 
Co(CN)2,  K4Co(CN)6;  cobaltic  compounds,  Co2O3,  Co(OH)3, 
Co2(SO4)3.i8H2O,  K3Co(NO2)6;  cobalt  amines,  Co(NH3)3Cl3.H26, 
Co(NH3)4Cl3.H2O,  Co(NH3)4Cl3.H2O,  Co(NH3)5Cl3,  etc. 

Experiment  I. — Properties  of  Cobalt. 

Examine  cobalt  metal  and  note  its  most  obvious  physical  properties  (  ?)• 
Scratch  the  metal  with  the  point  of  a  knife  blade.  Is  the  metal  hard? 
Does  it  tarnish  readily  in  the  air? 

Experiment  II. — Cobalt  Compounds. 

Examine  the  compounds  of  cobalt  (end  shelf).     Give  the  color  and 


364  EXPERIMENTAL  CHEMISTRY. 

empirical  formula  of  each  compound.     Record  the  valency  of  the  cobalt 
atom  in  each  compound.     What  is  the  color  of  the  cobalt  ion  ? 

Experiment  III. — Dehydration  of  Hydrated  Cobaltous  Chloride. 

Write  upon  a  sheet  of  your  note-book  with  a  solution  of  cobalt  chloride 
by  means  of  a  glass  rod.  Allow  it  to  dry,  then  warm  the  paper  very 
gently  by  holding  it  at  some  distance  from  a  gas  flame.  Result  ?  Breathe 
upon  the  paper  or  hold  it  for  an  instant  in  a  current  of  steam.  Explain 
the  changes  in  color.  Equations?  What  is  " sympathetic  ink "? 

Experiment  IV. — Reactions  of  Cobalt  Salts. 

Treat  separate  portions  of  a  solution  of  cobalt  chloride  with  the  fol- 
lowing reagents: 

(a)  Test  the  solution  with  a  borax  bead  in  the  oxidizing  (?)  and  re- 
ducing (?)  flames. 

(b)  Ammonium  hydroxide,  first  in  small  quantities  (?),  then  in  excess 
(?).     Equations? 

(c)  Sodium  hydroxide  solution,  first  in  small  quantities  (?),  then  in 
excess  (?).     Add  bromine  water  and  boil  (?).     Equations? 

(d)  Ammonium    sulphide    solution    (?).     Filter.     Try   the    effect   of 
dilute  hydrochloric  acid  upon  the  precipitate  (?).     Equations? 

(e)  Potassium  cyanide  solution,  first  in   small  quantities  (?),  then  in 
excess  ( ?) ;  add  sodium  hydroxide  solution  in  considerable  quantity,  then 
bromine  water*  until    the   color    of    bromine   persists;    warm  gently. 
Results  (?)    Equations? 

NICKEL,    Nl. 

At.  Wt.  58.7         Sp.  Gr.  8.8-9.0 

This  element  possesses  many  properties  in  common  with  cobalt. 
It  occurs  chiefly  in  combination  with  arsenic  as  niccolite,  NiAs,  and 
nickel  glance,  NiAsS.  It  is  now  manufactured  chiefly  from  garnierite, 
H4Ni2Mg2(SiO4)3.4H2O,  a  silicate  found  in  Australia.  The  crude 
nickel  may  be  obtained  from  granierite  by  reducing  the  ore  in  a  blast 
furnace  or  by  electrolysis.  Pure  nickel  is  usually  prepared  by  reducing 
the  oxide  with  carbon  at  a  high  temperature  or  by  reducing  the  oxide  in  a 
stream  of  hydrogen. 

Nickel  is  a  lustrous,  white  metal  (yellow-tinted),  very  hard  and  tena- 
cious, with  a  melting  point  at  1570°.  It  takes  a  very  high  polish.  It 
tarnishes  very  slowly,  even  in  moist  air.  On  account  of  its  resistance 
to  oxidation,  it  is  used  extensively  as  a  protective  covering  for  other 
metals  which  are  more  readily  oxidized,  such  as  iron,  etc.  The  process 
of  depositing  one  metal  upon  another  by  electrolysis  is  known  as  electro- 
plating. Nickel  forms  a  number  of  valuable  alloys.  German  silver  is  an 
alloy  of  copper,  nickel  and  zinc  (2  to  i  to  i).  Our  co-called  "nickel"  of 

*  Bromine  is  much  more  solublein  an  aqueous  solution  of  potassium  bromide  than 
in  water. 


TRANSITION    ELEMENTS.  365 

currency  contains  75  per  cent,  copper  and  25  per  cent,  nickel.  Nickel- 
steel  contains  from  4  to  15  per  cent,  of  nickel. 

Hydrochloric  acid  and  sulphuric  acid  attack  nickel  with  difficulty, 
but  nitric  acid  acts  upon  it  very  readily.  Concentrated  nitric  acid  renders 
the  metal  "passive." 

Nickel  forms  a  bivalent  ion,  nickelion,  Ni",  which  is  of  green  color. 
Nickel  can  form  a  higher  stage  of  oxidation,  for  example,  Ni2O3  and 
Ni(OH)3,  but  they  are  extremely  unstable  and  do  not  behave  as  base- 
forming  compounds. 

Nickel  can  also  form  complex  ions,  but  these  are  neither  so  numer- 
ous nor  so  stable  as  those  of  cobalt;  this  forms  the  essential  difference 
between  the  chemical  conduct  of  these  two  elements.  The  complex 
ions  of  nickel  containing  ammonia  also  differ  from  those  of  cobalt  not 
only  in  being  derived  from  bivalent  nickel,  but  also  in  being  very  unstable. 

Most  of  the  nickel  salts  are  green  when  hydrated  and  yellow  when 
dehydrated. 

The  following  formulae  show  the  composition  of  some  of  the  more 
important  compounds  of  nickel:  Nickelous  compounds,  NiO,  Ni(OH)2, 
NiCl2.6H2O,  NiSO4.6H2O,  Ni(NO3)2.6H2O,  NiCO3.6H2O,  NiS, 
Ni2(Fe(CN)6),  Ni3(Fe(Cn)6)2,  Ni(CN)2,  Ni(CN)2,  2KCN;  nickelic  com- 
pounds, Ni2O3,  Ni(OH)3. 

Experiment  I. — Properties  of  Nickel. 
Same  as  Experiment  I,  "  Cobalt." 
Experiment  II. — Nickel  Compounds. 
Same  as  Experiment  II,  "  Cobalt." 
Experiment  III. — Reactions  of  Nickel  Salts. 

Same  as  Experiment  IV,  "Cobalt."  A  solution  contains  both  the 
nickel  and  cobalt  ions.  How  may  the  cobalt  be  separated  from  the 
nickel  ? 

PROBLEMS. 

1.  It  was  found  that  1.586  grams  of  iron  formed  2.265  grams  of  ferric 
oxide.     Calculate  the  atomic  weight  of  iron. 

2.  How   many  grams  of   potassium   bichromate  will    be  required  to 
oxidize  10  grams  of  ferrous  sulphate? 

3.  How  many  grams  of  potassium  permanganate  will  be  required  to 
oxidize  10  grams  of  ferrous  sulphate? 

4.  What  weight  of  stannous  chloride  will  be  required,  theoretically, 
to  reduce  i  gram  of  ferric  chloride  to  ferrous  chloride  ? 


CHAPTER  XXXIV. 
TRANSITION  ELEMENTS 

Second  Long  Period.  Fourth  Long  Period. 

(Ruthenium,  Ru.  101.7)  (Osmium,  Os.  191.0) 

(Rhodium,  Rh.  103.0)  (Iridium,  Ir.     193.0) 

(Palladium,  Pd.   106.5)  Platinum,  Pt.    194.8 

These  rare  metals  constitute  two  separate  transitional  groups,  yet 
they  are  very  closely  related  to  one  another.  Again,  they  possess  certain 
properties  which  are  markedly  different.  The  first  three  have  atomic 
weights  which  are  close  to  one  hundred,  while  the  atomic  weights  of  the 
last  three  are  close  to  two  hundred.  The  specific  gravities  show  a  similar 
relation,  as  may  be  seen  from  the  following  table: 

I.  Ru,sp.gr.     12.26;     Rh,  sp.  gr.      12.10;     Pd,  sp.  gr.      11.9; 
II.    Os,  sp.  gr.     22.38;     Ir,  sp.  gr.        22.4;      Pt,  sp.  gr.       21.45. 

All  of  these  elements,  however,  resemble  platinum  more  or  less  closely. 
They  are  therefore  spoken  of  as  the  platinum  elements.  In  nature  they 
occur  associated  together  in  what  is  commonly  known  as  platinum  ore. 
This  ore,  which  is  sometimes  spoken  of  as  native  platinum,  is  found  in 
small  particles  and  nuggets  in  river  sand  and  alluvial  deposits.  The 
Urals  furnish  the  larger  portion  of  the  world's  supply;  smaller  quantities 
are  found  in  Australia,  California,  Borneo  and  Brazil.  The  ore  con- 
tains these  elements  in  the  metallic  state,  more  or  less  alloyed,  together 
with  small  quantities  of  iron,  copper  and  gold.  Platinum  constitutes  60 
to  85  per  cent,  of  the  ore.  Smaller  amounts  of  platinum  are  found  alloyed 
with  iridium. 

The  members  of  these  groups  are  lustrous,  white  metals,  unacted  upon 
by  air  at  ordinary  temperatures.  Osmium  burns  in  oxygen  when  highly 
heated,  forming  the  tetroxide,  OsO4.  The  other  five  metals  resist  oxi- 
dation at  any  temperature. 

Palladium  is  the  only  member  which  is  attacked  by  nitric  acid.  The 
other  members  are  not  acted  upon  by  ordinary  acids.  Aqua  regia  is 
without  action  upon  rhodium  and  iridium. 

Rhuthenium  is  a  hard,  brittle  metal,  fusing  at  about  2000°.  It  was 
discovered  by  Claus  in  1845.  Some  of  its  compounds  are  RuO,  Ru2O3, 
RuO2,  RuO4,  Ru(OH)3,  Ru(OH)4.3H2O,  RuCl2,  RuCl3,  RuCl4,  K2RuO4, 
KRoO4. 

Rhodium  is  a  malleable  metal,  fusing  at  2000°.  In  appearance,  it  re- 
sembles aluminum.  It  is  harder  than  platinum.  It  was  discovered  by 

366 


TRANSITION    ELEMENTS.  367 

Wollaston  in  1803.  Some  of  its  compounds  are,  RhO,  Rh2O3,  RhO2, 
Rh(OH)3,  Rh(OH)4,  RhCl3. 

Palladium  is  the  most  easily  fusible  of  these  metals,  melting  at  about 
1500°.  It  possesses  the  property  of  absorbing  (occluding)  large  volumes 
of  hydrogen  gas,  360  to  960  times  its  own  volume,  depending  upon  its 
state  of  aggregation  and  its  temperature.  When  heated  to  130°,  it 
surrenders  the  hydrogen.  It  has  been  supposed  for  a  long  time  that  the 
hydrogen  and  palladium  enter  into  a  definite  chemical  union  with  the 
formation  of  palladium  hydride,  Pd2H.  There  is  considerable  doubt  as 
to  whether  this  is  a  definite  chemical  compound.  The  later  explanation 
that  palladium-hydrogen  is  simply  a  "solid  solution"  in  which  hydrogen 
is  dissolved  in  palladium  does  not  satisfactorily  account  for  all  of  the 
observed  phenomena.  Palladium-hydrogen  is  a  powerful  reducing  agent 
owing  to  its  capability  of  releasing  hyrogen  in  a  condition  similar  to 
"nascent"  hydrogen  Some  of  its  compounds  are  Pd2O,  PdO,  PdCl2, 
PdCl4,  PdI2,  Pd(NO3)2,  PdSO4.2H2O,  H2PdCl6,  K2PdCl6,  (NH4)2PdCl6. 

Osmium  is  the  heaviest  known  elementary  substance.  It  fuses  at  2  500  °. 
An  alloy  of  iridium  and  osmium  which  is  very  hard  is  used  for  tipping  gold 
pens.  A  solution  of  osmium  tetroxide,  OsO4,  is  found  useful  in  hardening 
tissues  for  histological  purposes.  Some  of  its  more  important  compounds 
are  represented  by  the  following  formula?:  OsO,  Os2O3,  OsO2,  OsO4, 
Os(OH)2,  Os(OH)3,  OsCl2,  OsCl3,  OsCl4,  K2OsO4. 

Iridium  is  a  very  hard  metal  fusing  at  1950°.  It  alloys  with  platinum 
and  enhances  the  resistance  of  that  metal  to  the  action  of  acids.  It  is, 
therefore,  usually  present  in  platinum  utensils  designed  for  laboratory 
purposes.  The  following  formulae  represent  some  of  its  compounds: 
IrO,  Ir203,  Ir02,  Ir(OH)4,  IrCl2,  IrCl3,  IrCl4,  K3IrCl6.3H2O. 

Platinum  is  a  tough,  malleable,  ductile  metal  which  can  be  welded  at 
red  heat  It  fuses  at  1770°.  Its  temperature  coefficient  of  expansion  is 
about  the  same  as  that  of  glass.  Therefore,  if  it  is  desired  to  seal  an  elec- 
trical connection  through  glass,  a  platinum  wire  which  is  a  good  con- 
ductor is  the  most  convenient  means.  This  fact  is  utilized  in  construct- 
ing incandescent  electric  lights. 

Platinum  occludes  oxygen  and  hydrogen,  the  quantity  absorbed  de- 
pending upon  the  state  of  division  of  the  metal.  Platinum  black  (very 
finely  divided  platinum)  which  can  be  prepared  by  depositiong  it  from 
its  solutions  by  means  of  a  more  electro-positive  metal,  absorbs*  about 
300  times  its  own  volume  of  hydrogen  and  about  100  times  its  own 
volume  of  oxygen.  At  red  heat  the  gases  are  expelled  from  the  metal. 
The  value  of  platinum  as  a  "catalytic  agent"  is  due  to  its  capacity  to 
occlude  gases. 

Platinum  is  very  resistant  to  chemical  reagents,  and  upon  this  fact  its 
value  largely  depends.  It  is  not  attacked  by  the  ordinary  acids,  but  the 
free  chlorine  in  aqua  regia  converts  it  into  chloroplatinic.  acid,  H2PtCl6. 
It  is  not  acted  upon  by  the  fused  alkaline  carbonates,  but  it  interacts  with 
the  fused  alkalies,  giving  platinates.  The,  fused  alkaline  cyanides  also 

*  Hydrogen  is  "dissolved,"  but  oxygen  is  merely  concentrated  upon  its  surface. 


368  EXPERIMENTAL  CHEMISTRY. 

interact  with  it.  Platinum  must  not  be  heated  in  contact  with  carbon,  sili- 
con and  phosphorus,  as  they  unite  with  it  forming  compounds  which  are 
quite  brittle.  Lead  and  antimony  form  fusible  alloys  with  platinum, 
therefore  neither  these  metals,  nor  compounds  from  which  they  may  be 
liberated,  should  be  heated  in  platinum  vessels. 

Platinum  forms  compounds  in  which  it  is  the  positive  ion  showing  a 
valence  of  II  or  IV,  as  well  as  compounds  in  which  it  is  a  constitutent  of 
the  negative  ion.  However,  when  solutions  containing  ions,  in  which 
platinum  is  in  the  anion,  are  electrolyzed,  the  platinum  is  deposited  at  the 
anode  and  not  at  the  cathode.  The  following  list  contains  some  of  the 
more  important  compounds  of  platinum:  Platinous  compounds,  PtO, 
Pt(OH)2,  PtCl2,  PtS;  Platinic  compounds,  Pt(OH)4,  PtCl4,  PtS2;  chloro- 
platinic  acid,  H2PtCl6;  chloroplatinates,  K2PtCl6,  Na2PtCl6,  (NH4)2PtCl6; 
platinocyanides,  BaPt(CN)4.4H2O,  K2Pt(CN)4,3H2O. 

Experiment  I. — Properties  of  Platinum. 

(a)  Physical   properties.     Examine   specimens   of   platinum   foil   and 
wire.     Note  the  physical  properties  of  the  metal  (?).     Hold  a  piece  of 
platinum  wire  in  the  hottest  portions  of  the  Bunsen  flame  (?),  and  the 
flame  of  the  blast-lamp  (?).     What  is  "spongy  platinum"?     " Platinum 
black"? 

(b)  Chemical  properties.     Does  platinum  tarnish  readily  when  ex- 
posed to  the  action  of  the  air.     Procure  two  small  pieces  of  platinum 
scrap  from  the  assistant.     Place  them  in  separate  test  tubes.     Heat  the 
one  with  hydrochloric  acid  (?)  and  the  other  with  nitric  acid  (?).     Mix 
the  contents  of  both  test  tubes.     Result  ?     Equation  ? 

Experiment  II. — Platinum  as  a  "  Catalytic  Agent." 

Recall  or  repeat  those  experiments  in  which  platinum  acted  -as  a 
" catalytic  agent."  State  briefly  Ostwald's  tentative  explanation  of  the 
role  of  a  catalyzer.  (See  Chapter  IX,  Experiment  VII.) 

Experiment  III. — Platinum  Compounds. 

Examine  the  compounds  of  platinum  (end  shelf).  Give  the  color  and 
formula  of  each  compound.  What  is  the  color  of  the  platinum  ion? 

Experiment  IV. — Preparation  of  Platinic  Chloride. 

(a)  Scour  pieces  of  platinum  scrap  with  sea  sand,  then  wash  with  dis- 
tilled water,  and  boil  with  hydrochloric  acid.  Decant  the  liquid,  and 
again  wash  with  distilled  wrater.  Dissolve  i  gram  of  the  platinum  in  25 
cm.3  of  aqua  regia  in  a  covered  glass  dish.  After  the  "spurting"  has 
ceased  pour  off  the  supernatant  fluid  and  concentrate  it  in  a  glass  dish  by 
evaporation  on  the  steam  bath.  In  the  meantime  add  more  aqua  regia 
to  the  undissolved  platinum,  and  continue  the  above  operations  until  all 
of  the  platinum  is  dissolved.  After  the  combined  solutions  have  been 
evaporated  to  dryness,  moisten  with  a  little  hydrochloric  acid  and  take 
up  with  water,  or, 


TRANSITION    ELEMENTS.  369 

(b)  The  platinum  may  be  precipitated  from  the  concentrated  aqua- 
regia  solution  by  a  strong  ammonium  chloride  solution,  as  ammonium 
chloroplatinate,  (NH4)2PtCl6.  Filter,  ignite  the  precipitate  in  a  porcelain 
crucible.  Boil  the  residue  of  "spongy  platinum"  with  hydrochloric  acid, 
decant  the  liquid  and  dissolve  the  platinum  in  aqua  regia.  Evaporate  the 
solution  to  dryness  on  the  steam  bath,  moisten  the  residue  with  hydro- 
chloric acid  and  again  evaporate  just  to  dryness.  Take  up  the  product 
with  distilled  water.  Equations? 

Experiment  V. — Reactions  of  Platinum  Salts. 

(a)  Dilute  a  small  quantity  of  a  platinic  chloride  solution  with  2  or  3 
cm.3  of  water,  add  a  solution  of  hydrogen  sulphide,  and  warm  slightly  ( ?). 
Is  the  precipitate  soluble  in  excess  of  potassium  hydroxide  ?     In  ammo- 
nium polysulphide  ? 

(b)  Reduction  of  platinum  salts.     "Platinum  black".*     To  a  solution 
of  platinic  chloride  add  an  excess  of  sodium  carbonate  solution  and  a  little 
grape  sugar,  then  boil  the  mixture.     Carbon  dioxide  is  evolved,  and  a 
black  powder  ("platinum   black")  is   formed   slowly.     The    solutions 
should  be  dilute.     Wash  the  powder  successively  with  dilute  alcohol, 
hydrochloric  acid,  potassium  hydroxide  solution,  distilled  water,  and  then 
dry  it  by  applying  a  gentle  heat. 

(c)  Recall  the  interactions  of  solutions  of  platinic  chloride  with  the 
alkaline  chlorides.     (See  "Sodium,"  "Potassium,"  and  "Ammonium.") 

For  a  discussion  of  the  properties  of  praseodymium,  neodymium,  sama- 
rium, gadolinium,  terbium,  erbium,  thulium,  ytterbium,  etc.,  see 
reference  texts. 

*"  Platinum  black"  is  a  more  effective  catalyzer  than  "spongy  platinum." 


CHAPTER  XXXV. 

RELATIONS  WITHIN  THE  GROUPS  or  THE  PERIODIC  CLASSIFICATION. 

TABLE. 
Grouping  of  the  Metals  (Cations)  for  Purposes  of  Analysis. 

METALS  (CATIONS) 

_  !  __ 


II. 


Div.  A. 


Div.  B. 


III. 


IV. 


V. 


Metals  which  are  precipitated 
from  solutions  by 

HC1 
Silver 

Mercury  (ous) 
Lead 


I 

Metals  which  are  not   precipitated 
from  solutions  by 

HC1 


Sulphides  soluble  in  ammonium  polysulphide. 


Metals  which  are  precipitated 
from    solutions    containing 
hydrochloric  acid  by 
H2S 

Mercury  (ic) 

Copper 

Cadmium 

Bismuth 

Arsenic 

Antimony 

Tin 

Gold 

Platinum 

Metals  precipitated  from  ammoni- 
acal  solutions  containing  ammo- 
nium chloride  by 
(NH4)2S 

Cobalt 

Nickel 

Iron  Chromium 

Aluminum 

Zinc 

Manganese 


I 

Metals  which  are  not  precipitated 
from  solutions  containing  hy- 
drochloric  acid  by 

H2S 


Metals  not  precipitated  from  ammoni- 
acal  solutions  containing  ammonium 
chloride  by 

(NH4)2S 


I 

Metals  which  are  precipitated 
from  solutions  by 

(NH4)2C03 
Barium 
Calcium 
Strontium 


Metals  which  are  not  precipitated 
from  solutions  by 

(NH4)2C03 


Metal  which  is  precipitated 
from  solutions  by 
NaNH4HPO4 

Magnesium 


VI. 


Metals  which  are  not  precipitated 
from  solutions  by 

NaNH4HPO4 
Sodium 
Potassium 
Ammonium 


370 


EXPERIMENTAL    CHEMISTRY. 


371 


TABLE. 

GROUPING  or  NON-METALLIC  RADICALS  (ANIONS)  FOR  PURPOSES  or  ANALYSIS. 


ACID-RADICALS  (ANIONS) 

I 


Radicals  which  are  precipitated 
from  neutral  solutions  bv 


BaCl2 


J 

Radicals  which  are  not  precipitated 
from  neutral  solutions  by 

BaCl2 


Sulphuric,  H3SO4 
Hydrofluosilic,  H2SiF6 

Boric,  H3BO3 
Carbonic,  H2CO3 
Citric,  H3(C6H567) 
Chromic,  H2CrO4 
Hydrofluoric,  HF 
lodic,  HIO3 
Oxalic,  H2(C2O4) 
Phosphoric,  H3PO4 
Sulphurous,  H2SO3 
Silicic,  H2SiO3 
Tartaric,  H2(C4H4O6) 
Thiosulphuric,  H2S2O3 


Barium  salts  insoluble  in  dilute 
hydrochloric  acid. 


Barium  salts  soluble  in  dilute 
hydrochloric  acid. 


I 

II.     Radicals  precipitated  from  solutions 
acidified  with  nitric  acid  by 


AgN03 


Radicals  which  are  not  precipitated 
from  solutions  acidified  with  nitric 
acid  by 

AgN03 


Ferricyanic,  H3Fe(CN)6 
Ferrocyanic,  H4Fe(CN)6 
Hydrochloric,  HC1 
Hydrobromic,  HBr 
Hydriodic,  HI 
Hydrocyanic,  HCN 
Hydrosulphuric,  H2S 
Thiocyanic,  HCNS 


Acetic,  H(C2H3O2) 
Chloric,  HC1O3 
Perchloric,  HC1O4 
Cyanic,  HCNO 
Formic,  H(CHO2) 
Nitric,  HNO3 
Nitrous,  HNO2 


APPENDIX  I. 


PRELIMINARY  EXERCISES. 
APPARATUS. 

Since  chemistry  is  pre-eminently  an  experimental  science,  it  is  in  order 
to  begin  work  by  examining  various  pieces  of  apparatus  and  acquainting 
one's  self  with  their  manipulation. 

THE    BUNSEN    BURNER. 

Carefully  examine  the  Bunsen  burner  (Fig.  49).  Take  it  apart;  assemble 
it,  noting  relation  of  parts.  Connect  it  with  the  gas  supply.  Light  the 
burner  by  turning  on  the  gas,  and  then  holding  a  lighted  match  near  the 
side  of  the  burner  a  short  distance  below  the  top.  Notice  the  effect  of 
opening  and  closing  the  holes  at  the  bottom  of  the  tube.  Explain.  Ad- 
just the  opening  until  the  luminous  region  has 
just  disappeared  and  the  flame  is  noiseless: 
this  is  the  "Bunsen  flame."  A  flame  6  cm.  to 
9  cm.  high  is  adapted  to  most  work.  If  soot  is 
deposited  upon  the  object  being  heated,  open 
the  holes  at  the  foot  at  the  tube,  but  not  so  far 
as  to  produce  a  noisy  flame. 

Now  study  the  structure  of  the  flame.  De- 
termine which  parts  are  relatively  hotter  and 
which  cooler.  Introduce  quickly  into  the 
center  of  the  flame  (the  darker  portion),  about 
one-half  of  a  centimeter  above  the  burner,  the 
head  of  a  match.  Results  ?  Can  you  insert  a 
match  as  per  above  directions  and  withdraw  it 
without  igniting  same? 

Hold  a  piece  of  platinum  wire  across  the 
flame  in  various  places.  Note  the  color  of  the 
wire.  Bring  quickly  and  horizontally  (to  desk 
top)  a  piece  of  heavy  pasteboard  into  the  flame, 
about  2  cm.  above  the  top  of  the  tube;  hold  it  quietly  for  several  seconds, 
then  remove  quickly.  Do  you  find  a  brown  (scorched)  circle  ?  Explain. 

To  secure  the  greatest  heating  effect,  place  object  just  above  apex  of 
the  dark  inner  cone  of  unburned  gas.  Why  is  there  a  cone  of  unburned 
gas  ?  Why  the  shape  ?  If  oxygen  is  necessary  for  the  combustion  of  the 
gas,  where  might  the  gas  escape  for  a  time  unburned? 

373 


FIG.  49. — Bunsen  Burner. 


374  EXPERIMENTAL  CHEMISTRY. 

Half  fill  a  test  tube  with  water.  Be  sure  that  the  outside  of  the  tube  is 
dry.  By  means  of  test  tube  holders,  introduce  the  tube  into  the  hottest 
region  of  the  flame,  inclining  the  tube  at  an  angle  of  about  45°  to  the  top 
of  the  desk.  Heat  only  that  portion  of  the  tube  containing  the  liquid: 
if  the  flame  strikes  the  tube  above  the  liquid,  the  tube  may  crack. 
Keep  the  liquid  in  the  tube  slightly  agitated  by  a  short,  quick  move- 
ment of  the  hand. 

Above  directions  should  be  observed  in  all  cases  where  liquids  are 
heated  in  test  tubes.  A  "wing-top"  attachment  gives  a  broad  flame 
of  very  much  use  to  the  glass-blower. 

Note. — If  gas  is  not  available  in  the  laboratory  and  alcohol  lamps  are 
used,  perform  as  many  as  possible  of  above  experiments. 

THE    BLAST-LAMP. 

Where  a  much  higher  temperature  is  required  than  can  be  secured  by 
means  of  the  Bunsen  burner,  the  blast-lamp  is  used.  The  size  of  the 
flame  can  be  altered  by  proper  manipulation.  (Instructions  from  assistant.) 

MANIPULATION    OF    GLASS. 

I. — To  Cut  Glass  Tubing. 

(a)  Lay  the  tubing  on  a  flat  surface;  make  a  file-mark  on  it  at  right 
angles  to  the  length;  take  the  tube  in  the  hands,  placing  the  two  thumbs 
opposite  the  scratch  and  the  fingers  on  either  side  of  the  scratch;  now  push 
gently  with  the  thumbs  and  at  the  same  time  pull  the  hands  apart,  the 
tubing  usually  breaks  squarely  at  the  scratch.     " Fire-polish"  the  ends 
of  tube  by  turning  them  slowly  in  the  Bunsen  flame. 

(b)  To  break  large  tubing  or  cut  off  bottoms  of  bottles,  etc.,  encircle 
tube  with  wire  or  make  an  ink-mark  to  trace  the  path  of  the  desired 
break  or  cut,  then  a  file-mark  is  made  upon  the  surface  on  mark;  a  steel 
file  handle,  or  better,  a  glass  rod  heated  in  the  blast-lamp  flame  until  it  is 
red  hot,  when  it  is  at  once  pressed  against  the  scratch  until  the  glass 
begins  to  crack.     The  fracture  can  usually  be  led  in  any  direction  by  keep- 
ing the  hot  glass  rod  in  front  of  it.     Heat  rod  frequently  to  keep  it  red  hot. 

II. — Grinding  Glass. 

(a)  Rough  edges  of  tubes  or  bottles  may  be  ground  down  to  bell-jar 
effect  by  spreading  emery  paste  upon  a  smooth  flat  surface  and  rubbing 
broken  edges  upon  it. 

(b)  Glass  stoppers  may  be  ground  into  necks  of  flasks,  etc.,  by  covering 
stoppers  and  inside  of  neck  with  emery  paste,  and  then  imparting  a  gentle 
pressure  to  stopper  while  twisting  it  into  place. 

III. — Cutting  and  Perforation  Glass  Plates. 

(a)  The  plate  of  glass  is  laid  upon  a  flat  surface;  "glass  cutters" 
are  used  to  make  the  scratch  where  the  break  is  desired.  "  Glass  cutters" 
are  usually  made  with  a  rotating  wheel  of  steel  or  a  diamond  point. 


APPENDIX    I.  375 

(b)  Holes  can  be  made  in  a  glass  plate  by  the  aid  of  a  broken  end  of  a 
round  file  kept  wet  with  a  solution  of  camphor  in  oil  of  turpentine. 

IV. — To  make  Stirring  Rods. 

Cut  off  a  piece  of  glass  tubing  18  cm.  to  20  cm.  long  and  6  mm.  in  di- 
ameter. Hold  the  ends  of  the  tube  successively  in  a  Bunsen  flame,  rotating 
the  tube  constantly  until  the  open  ends  are  sealed.  Glass  rods  may  be 
used  instead. 

V. — To  Bend  Glass  Tubing. 

A  flat  Bunsen  flame,  produced  by  a  "wing-top"  or  a  "fish-tail"  attach- 
ment, is  used  for  bending  glass.  Take  the  tube  in  both  hands  and  hold 
that  portion  which  is  to  be  bent  lengthwise  to  the  flame  and  just  above  the 
flame  until  it  is  warmed;  then  place  it  in  the  flame,  constantly  rotating  on 
the  long  axis,  between  thumb  and  fingers,  until  glass  becomes  fairly  soft; 
remove  it  from  flame  and  quickly  bend  it  into  the  desired  form.  It  is 
well  to  anneal  the  glass  at  the  bend  by  "smoking"  it.  This  maybe  ac- 
complished by  closing  the  holes  at  the  base  of  the  burner,  thus  producing 
a  smoky  flame.  The  bent  portion  should  not  be  permitted  to  touch  cold 
objects  until  it  has  cooled. 

Using  ordinary  glass  tubing,  the  student  should  make  various  styles  of 
bends  like  models  shown  by  the  assistant.  Always  "fire-polish"  the 
edges  of  glass  tubing. 

Note. — To  increase  the  internal  diameter  of  tubes  for  insertion  of  corks, 
etc.,  soften  the  tube  in  the  flame,  and  insert  a  conical  piece  of  charcoal 
by  gentle  pressure  until  tube  spreads  into  desired  shape. 

VI. — Joining  Tubes  and  Glass  Blowing. 

See  Ostwald's  "Physico-Chemical  Measurements,"  pages  66-72,  also 
Shenstone's  "  Methods  of  Glass  Blowing."  The  student  will  find  in  said 
references  brief  but  excellent  discussions  of  the  subject. 

VII. — Sealing  Platinum  Wires  into  Glass  Tubes. 

Soften  a  glass  tube;  draw  it  out;  cut  off  short,  and  by  heating  cause  the 
end  to  fall  nearly  together,  or  until  the  wire  can  just  be  pushed  into  the 
opening;  heat  until  glass  closes  around  the  wire. 

More  certain  results  are  obtained  if  the  platinum  wire  receives  a  drop 
of  melted  enamel  (tough  lead-glass)  at  the  proper  place  and  is  pushed 
through  opening  in  tube  until  hole  is  closed  by  the  enamel.  The  enamel 
unites  well  with  platinum  and  also  with  ordinary  glass. 

PERFORATION   OF    STOPPERS. 

A  set  of  cork  borers  may  be  secured  from  the  assistant.  (T.O.) 
Hold  the  cork  in  the  hand  and  bore  from  the  narrow  end.  Avoid  great 
pressure  on  the  cork  borer. 


376 


EXPERIMENTAL  CHEMISTRY. 


In  perforating  rubber  corks,  the  borer  cuts  more  easily  if  it  is  dipped 
frequently  into  a  solution  of  caustic  soda.  A  round  file  may  be  used  to 
smooth  the  perforation. 

TREATMENT    OF    RUBBER    CORKS    AND    TUBING. 

Where  rubber  is  to  be  used  in  quantitative  experiments,  it  should  be 
boiled  in  dilute  sodium  hydroxide  solution,  rinsed  with  water,  then  boiled  in 
dilute  hydrochloric  acid,  and  finally  washed  with  water.  This  operation 
removes  impurities  which  frequently  introduce  errors. 


FIG.  50.— Water  Bottle. 


FIG.  51. — Mohr  Burettes. 


CONSTRUCTION  OF  PARTS  AND  THE  ASSEMBLING  OF  A  WASH  BOTTLE. 

Select  proper  material  and  construct  a  wash  bottle  (Fig.  50)  like  model 
in  laboratory.  The  necessary  material  will  be  found  in  the  drawer. 

MEASURING    INSTRUMENTS. 

(a)  Measures  of  volume. 

Assemble  the  Mohr  (Fig.  51) ;  or  Geissler  burette,  and  clamp  it  in  a  verti- 
cal position.  Observe  the  model.  Fill  the  burette  with  distilled  water ;  avoid 
air  bubbles  in  stop  cock;  run  out  water  into  a  beaker  until  the  lower  side 
of  meniscus  (Fig.  54)  (curved  surface)  stands  at  40.85  c.c.;  estimate  to 


APPENDIX    I. 


377 


tenths  of  a  division;  record  this  reading  in  your  note-book;  now  run  out 
water  into  a  "graduate  "  until  the  latter  is  about  half  filled;  take  the  read- 
ing on  the  "graduate";  how  many  cm.*?  Take  the  burette  reading; 
how  many  cm.3  did  you  run  out  ?  Compare  readings.  Which  gives  the 
more  accurate  measurement — burette  or  graduate  ?  Why  ? 

By  means  of  a  calibrated  pipette  (Fig.  55),  introduce  its  nominal  volume 
of  water  into  a  "graduate."  Compare  readings.  Which  is  the  more 
accurate  ?  Why  ? 

By  use  of  the  graduate,  determine  the  approximate  volumes  of  water 
which  test  tubes,  beakers  and  flasks  will  hold.  Make  a  record  of  data. 


FIG.  52. — Burette 

Operated  with 

Pinch-cock. 


FIG.  53. — Burettes  with  Glass 
Stop-cocks. 


FIG.  54. — Meniscus. 


By  aid  of  burette,  determine  the  volume  occupied  by  a  "drop"  of  water 
as  delivered  by  your  burette.  Hint. — Run  out  ten  to  twenty  drops.  Make 
a  record  of  exercise. 

(b)  Measure  of  temperature. 

The  mercury-glass  therometer  is  ordinarily  used  in  the  laboratory. 
Note  the  possible  maximum  and  minimum  readings.  Always  observe 
this  before  using  instrument.  Some  thermometers  are  made  to  read  to 
only  -f  60°  C.;  such  a  thermometer  plunged  into  water  boiling  at  any 
temperature  above  that  registered  on  scale  would  undoubtedly  be  ruined. 
Never  plunge  the  instrument  into  relative  extremes  of  temperature. 

Suspend  the  thermometer  in  a  beaker  of  water  which  has  presumably 
acquired  the  temperature  of  the  laboratory.  Remember  the  heat  of  the 


EXPERIMENTAL  CHEMISTRY. 


25C 


hand  will  affect  the  reading.  It  requires  time  for  a  thermometer  to  come 
to  the  temperature  of  a  new  environment — the  thermometer  is  said  to 
"lag."  Gently  tap  the  thermometer  with  finger  before  reading — 
this  is  to  overcome  "stiction."  What  is  the  temperature  of  the  water  in 
beaker  on  the  Centigrade  scale?  Fahrenheit  scale?  Reaumur  scale? 
Make  a  record  of  experiment. 

The  Beckmann  thermometer  for  deter- 
mining accurately  small  changes  in  temper- 
ature is  frequently  used.  Consult  the  in- 
structor with  reference  to  its  manipulation. 
(c)  Measures  of  weight. 
The  equal  arm  lever  balance  is  the  in- 
strument most  frequently  used  in  the 
chemical  laboratory  for  the  determination 
of  weight.  The  principle  of  this  balance  is 
embodied  in  two  useful  forms — the  "plat- 
form" or  "trip"  balance  and  the  so-called 
"analytical"  balances  (Fig.  56).  Both 
kinds  will  be  found  in  the  laboratory.  Ex- 
amine them,  the  latter  under  the  supervision 
of  the  assistant.  The  former  is  used  in 
making  those  weighings  where  only  an  ap- 
proximate accuracy  is  demanded;  the  latter 
are  much  more  sensitive  and  accurate,  and 
should  be  used  only  when  the  experiment 
is  marked  "Quant." 

The  balance  and  weights  must  be  handled 
with  care.  Solids  to  be  weighed  must  be 
placed  first  upon  a  piece  of  paper  or  a  watch 
glass,  never  directly  upon  the  pan.  All 
objects  must  be  perfectly  clean  and  dry. 
The  weights  must  be  handled  with  forceps 
— not  with  the  moist  hand. 

The  general  procedure  when  weighing 
with  platform  balances  is  as  follows:  find 
zero  point  of  balance  by  allowing  the  beam 
to  swing,  noting  whether  the  pointer  makes 
equal  excursions  on  either  side  of  the  zeVo 
mark;  if  it  does  not,  correct  defects  by 
placing  pieces  of  paper,  etc.,  on  the  proper 
pan.  This  is  called  "counterpoising"  a 

-pic.  55. — Pipettes.  balance.     In  the  future,  when  weighing,  do 

not  wait  for  pointer  to  come  to  rest  at  the 

zero  point,  simply  add  to  or  subtract  weights  from  the  proper  pan  until 
the  vibrations  on  either  side  of  zero  point  are  of  equal  amplitude. 

The  assistant  will  instruct  you  as  to  the  proper  method  of  weighing  with 
the  "analytical"  balances. 


APPENDIX    I. 


379 


Note. — To  avoid  one  of  the  commonest  errors  in  weighing,  count  the 
values  of  the  vacant  places  in  the  set,  then  check  by  counting  the  weights 
in  the  pan.  Record  on  paper  the  value  of  the  weights  as  you  take  them 
from  the  pan. 


CALIBRATING    BY    WEIGHING. 


Make  a  mark  or  paste  a  piece  of  gummed  paper  on  the  lower  portion 
of  the  neck  of  a  50  cm. 3  Erlenmeyer  flask.  Weigh  it  first  on  the  platform 
balances,  then  on  the  analytical  balances.  (Instructions.)  Place  flask 
on  left  pan.  Record  weights.  Now  fill  flask  to  mark  with  water  at 


© 


© 


FIG.  56. — Chemical  Balance. 

20°  C.  and  weigh  as  before.  Record  weights.  If  a  cm.3  of  water  at 
20°  C.  weighs  .998  gram,  what  is  the  volume  of  flask  when  filled  to 
mark?  Which  is  buoyed  up  the  most  by  air — flask  and  contents  or 
weights?  Is  this  a  source  of  error? 


PROBLEMS. 

1.  Express  10  cm.  in  millimeters,  decimeters  and  kilometers. 

2.  Convert   115  in.  into  centimeters;  98  mm.  into  inches;  2  yd.  into 
meters;  2  ft.  into  centimeters;  760  mm.  into  inches. 

3.  How  many  liters  in  10  quarts?     Quarts  in  4  liters? 

4.  How  many  liters  will  a  pneumatic  trough  hold,  the  dimensions  of 
which  are  15  in.x2o  in.x25  in.? 

5.  Add  9  grm.,  468  mg.,  7  dg.,  and  5  eg.,  and  express  the  sum  in  grams. 

6.  10  kg.  are  equivalent  to  how  many  pounds?     Convert  10  oz.  into 
grams. 


380  EXPERIMENTAL    CHEMISTRY. 

7.  How  many  grams  in   134.76  dg.  ?     In  17,589  eg.?     In  5.95  mg.  ? 

8.  How  many  cubic  centimeters  (cm.3)  in  .5!.?     In  .75!.?  In  95  dm.3 

9.  Convert  32°  Fahr.  into  a  Centigrade  reading.     A  Reaumur  read- 
ing. 

10.  Repeat  9,  using  212°  Fahr. 

11.  A  thermometer  bearing  a  Fahrenheit  scale  registers  a  temperature 
of  72°.     What  would  be  the  equivalent  reading  on  the  Centigrade  scale? 
On  the  Reaumur  Scale? 

12.  Convert  20°  C.  into   an   equivalent  reading  on  the  Fahr.  scale. 
On   the  Reaumur  scale. 

13.  The  sp.  gr.  of  concentrated  sulphuric  acid  is  1.84,  of  nitric  acid 
1.4,  and  of  hydrochloric  acid  1.2.     Calculate  the  weight  of  a  liter  of  each 
acid. 

14.  How  many  cm.3  in  10  grm.  of  each  of  the  above  acids? 

15.  Alcohol   (ethyl)   at   15°  C.  has  a  density  of  .7937.     What  is  the 
weight  of  one  1.  at  15°  C.     How  many  cm.3  in  20  grm.? 

1 6.  An  empty  flask  has  a  weight  of  96.75  grm.;  when  filled  with  water 
at  20°  C.  it  has  a  weight  of  596.30  grm.     What  is  the  capacity  of  the 
flask  at  20°  C.  ?     Hint. — calculate  the  number  of  grams  of  water  which 
the  flask  holds  and  multiply  by  the  "correction  factor,"  1.0028.     This 
factor  corrects  for  the  buoyant  effect  of  the  air  in  weighing,  the  density  of 
water  (.99823)  at  20°  C.,  and  the  cubical  coefficient  of  expansion  of  glass 
(.000025).     The  "correction  factor"  for  19°  C.  is  1.0027.  See    Appendix. 
In  standardizing  volumetric   apparatus  the  "correction  factor"  should 
always  be  used. 


APPENDIX  II. 


NOTE. — The  data  contained  in  the  following  tables  have  been  selected 
and  arranged  from  various  sources.  Effort  has  been  made  to  incorporate 
the  results  of  the  most  recent  investigations.  In  many  instances  it  has 
been  possible  to  give  only  approximate  values. 

The  tables  which  have  been  most  frequently  consulted  are:  Landolt 
und  Bornstein,  Physikalisch-chemische  Tabellen;  Ostwald,  Physico- 
chemical  Measurements;  Kohlrausch,  Praktische  Physik;  Buchka, 
Physikalisch-chemische  Tabellen  der  unorganischen  Chemie;  Watson, 
Practical  Physics;  Miller,  Laboratory  Physics;  Smithsonian  Physical 
Tables;  Comey,  Dictionary  of  Solubilities;  Seidell,  Solubilities  of  Inorga- 
nic and  Organic  Substances;  J.  Thomsen,  Thermo-chemische  Unter- 
suchungen. 

TABLE  I. 
Metric  Measures  with  English  Equivalents. 

(a)  Measures  of  Length. 

i  Millimeter,  mm.  =  .0393  inches, 

i   Centimeter,  cm.  =10  mm.    =  .3937  inches, 

i  Decimeter,  dm.  =10  cm.     =  3.9371  inches, 

i  Meter,  m.  =10  dm.     =39.3708  inches, 

i  Meter,  m.  =  3  .  2809  feet, 

i  Meter,  m.  =  1.0936  yards, 

i  Kilometer,  Km.  =1000  m.     =  .6114  miles. 

(b)  Measures  of  Volume. 

i  Cubic  centimeter,  cm.3  =  .001  1.  =  .06103  cu-  inches. 

i  Cubic  decimeter  (liter),  i.  =  1000  cm.3  =  61.227  cu-  inches. 

i  Liter  =  2.1134  pints  U.S.  =  1.067  quarts  U.  S.=  0.26417  gallons  U.  S. 

i  Cubic  meter  =  1000  1.  =  35.317  cu.  ft. 

(c)  Measures  of  Weight. 

Milligram,  mg.  —  .001  gram  =      .0154  grains. 

Centrigram,  eg.  =  .01     gram  =      .1543  grains. 

Decigram,  dg.  =  .1      gram  =   1.5432  grains. 

Gram,  gm.  =  15.432  grains  =      .0352  av.  ounces. 

Gram  =     .03216  troy  ounces. 

Kilogram  (kilo)  =  1000  grams  =   2 . 2055  av.  Ibs. 

Kilogram  ==   2 . 6803  troy  Ibs. 

38* 


2  EXPERIMENTAL    CHEMISTRY. 

TABLE  II. 
English  Measures  with  Metric  Equivalents. 

(a)  Measures  of  Length. 

i  Inch   =25.399  mm-  =  2-539  centimeters. 

i  Foot   =12  inches  =  .3048  meter. 

i   Yard  =3  feet  =  .9144  meter. 

i  Mile   =1700  yards  =  5280  feet. 

i  Mile   =  i . 609  kilometers  =  1609.3  meters. 

(b)  Area.  (c)  Contents. 

i  Square  inch  =  6.4514  cm.2       i   Cubic  inch  =  16.386  cm.3 

i  Square  foot    =  929.01  cm.2       i   Cubic  foot  =  28.316  liters, 

i  Square  yard  =  8361.1  cm.2      i  Cubic  yard-  =  764.52  liters, 

i  Square  yard  =  .83661  m.2        i   Cubic  yard  =  about  .76  m.3 

(d)  Measures  oj  Volume. 

i  Pint    =  .47317  liter  =     473.11  cubic  centimeters, 

i  Quart  =  .94634  liter  =     946.22  cubic  centimeters, 

i   Gallon(U.S.)  =  3.785  liters  =     231        cubic  inches. 

(e)  Apothecaries'  Fluid  Measure,  U.S. 

i  Minim  =  a  drop  (approx)  .0616  cubic  centimeters, 

i  Fluid  dram  =  60  minims  3  . 6965  cubic  centimeters. 

i  Fluid  ounce  =  8  fluid  drams     =       29.572    cubic  centimeters, 
i  Pint  =  16  fluid  ounces  =     473.11       cubic  centimeters. 

(/)     Measures  of  Weight. 

I.     (Avoirdupois.) 

i   Gram  =  64.773  milligrams  =  0.0648  grams, 
i  Dram  =  27.34  grains  =  1.772  grams, 
i   Ounce  =  1 6  drams  =  437.5  grains  =  28.349  grams. 
i  Pound  =  16  ounces  =  7000  grains  =  453.59  grams. 
i   Short  ton  =  2000  pounds  =  907.17  kilograms, 
i  Long  ton  =  2240  pounds  =  1015.03  kilograms. 

II.     (Troy.) 

i   Grain  =  64.773  milligrams  =  .0648  grams. 
i  Pennyweight  =  24  grains  =  1.555  grams, 
i   Ounce  =  20  pennyweights  =  480  grains  =  31.103  grams, 
i  Pound  =  12  ounces  =  5760  grains  =  373.242  grams. 

III.     (Pharmacy  or  Apothecaries\) 
i   Grain  =  64.773  milligrams  =  .0648  grams, 
i  Scruple  =  20  grains  =  i .  296  grams, 
i  Dram  =  3  scruples  =  60  grains  =  3 . 888  grams, 
i   Ounce  =  8  drams  =  480  grains  =  31 . 1035  grams, 
i  Pound  =  12  ounces  =  373.248  grams. 


APPENDIX    II. 

TABLE  III. 


383 


Conversion  of  Thermometric  Readings. 

The  temperatures  mentioned  in  this  book  are  expressed  in  terms  of 
the  Centigrade  scale.     There  are  three  scales  now  in  general  use.     They 


are: 


I.     Fahrenheit—  F.    Water  freezes  at  32°,  boils  at  212°. 
II.     Centigrade  —  C.    Water  freezes  at    o°,  boils  at  100°. 
III.     Reaumur   —  R.    Water  freezes  at    o°,  boils  at  80°. 


To  convert:     F.  to  C. 


=  C°. 


5(F.°— 32°) 
or — =  C°. 

F.  to  R.     4(F-°~32°)  =  R> 


C.  to  F.     (C.  x  1.8)  +  32°  =  F.°. 


or 


=  F. 


R.toF. 


TABLE  IV. 

Corrections  to  Reduce  Readings  on    Mercury-in-Glass  Thermometers 
to  the  Normal  Hydrogen  Scale  or  Air  Thermometer. 

(For  Jena  Normal  Glass,  16'")* 


Reading 
Correction 

o°C. 
o°.ooo 

10°  C. 

—  °-°55 

20°  C. 

—  0.090 

30°  c. 

—  0.109 

40°  C. 
—  0.116 

—  0.109 

Reading 
Correction 

50°  c. 

—  o°.io9 

60  C. 

—  0.096 

70°  C. 
—  0.076 

80°  C. 
—  °-°53 

90°  C. 

-  —  0.027 

100°  C. 

0.000 

*From  tables  published  by  Griitzmacher  in  Wied.  Annalen  (1899)  p.  769. 


EXPERIMENTAL    CHEMISTRY. 


TABLE  V.     DENSITY  OF  WATER.* 

(Ax  THE  TEMPERATURE  "/"  ON  THE  NORMAL  HYDROGEN  SCALE) 


8 

1^ 

Q 

Tenths. 

o 

i 

2 

3 

4 

5 

6     7 

8 

9 

0° 

0-99987 

87 

88 

89 

89 

90 

90    91    92 

92 

i 

2 

3 
4 

93 
97 
99 

I  '  00000 

93 
97 

99 

oo 

94 

97 
99 

00 

94 

98 

00 
00 

95 
98 

00 
00 

95 
99 

00 

oo 

95    96 
99    99 

00      00 

oo  ;  oo 

96 
99 

00 
00 

97 
99 

00 

00 

5 

0-99999 

99 

99 

99    98 

98 

98    98    97    97 

6 

8 
9 

97 
93 
88 
81 

97 
93 
87 
80 

96 
92 
86 
79 

96 
92 
86 

79 

95 
9  1 

78 

95 
90 
84 
77 

95    94 
90    89 
84    83 
76    75 

94 
89 
82 

74 

93 
88 
82 
74 

64 

10 

73 

72 

7i 

70 

69 

68 

67    66 

65 

ii 

12 
13 
14 

63 

53 
40 

27 

62 
5i 
39 
26 

61 

5° 
38 
24 

60 
49 
37 
23 

59 
48 

35 

22 

58 
47 
34 
20 

57 
45 
33 
19 

56 

44 
3i 
17 

55 
43 

3° 
16 

54 
42 
29 
14 

15 

13 

ii 

10 

08 

06 

°5 

°3 

02 

00 

99 

16 
17 
18 
19 

0-99897 
80 
62 
43 

95 
78 
60 
4i 

94 
77 
59 
39 

92 

75 
57 
37 

9° 
73 

55 
35 

89 
71 
53 
33 

87 

70 

5i 
3i 

85 

68 

49 
29 

84 
66 

47 

27 

82 
64 
45 
25 

20 

23 

21 

i9 

17 

15 

13 

ii 

08 

06    04 

21 
22 

23 
24 

02 
0-99780 

56 

32 

00 

77 
54 
30 

98 
75 
52 
27 

95 
73 
49 
25 

93 
7i 

47 

22 

9i 
68 

45 

20 

89 
66 

42 
i7 

86 
64 
40 
15 

84 
61 

37 

12 

82 
59 
35 
10 

25 

07 

°5 

02 

99 

97 

94 

92    89    86 

8j 

26 

11 

29 

0-99681 

11 

o'  99597 

78 
5i 
23 
94 

76 
48 
2O 
91 

7l 
46 

17 

88 

70 

43 
15 
85 

67 
40 

12 

82 

65 
37 
09 

79 

62    59 

34    32 
06    03 

76    73 

57 
29 

00 

70 

30 

67 

64 

61 

58 

55 

S2 

49 

46 

43 

40 

31 
32 

33 
34 

37 
05 
o'99473 
40 

34 

02 
70 
36 

3£ 

99 
66 

33 

27 
96 
63 
3° 

24 

?2 
60 

26 

21 

89 
56 
23 

18 
86 
53 

20 

15 
83 

5° 
16 

12 

79 
46 

13 

08 
76 

43 
09 

35 

06 

02 

99 

95 

92 

89 

85 

82 

78 

75 

*Watson. 


APPENDIX   II. 


385 


TABLE  VI. 

VOLUME  OF  ONE  GRAM  OF  WATER.* 

(Thiesen,  Scheel,  and  Diesselhorst,  in  Wiss.  Abh.  d.  Phys.  Tech.  Reichsanstalt,  Vol. 
II,  1895;  Vol.  Ill,  1900;  and  Vol.  IV,  1904.) 


Tem- 
pera- 
ture. 

0 

i 

2 

3 

4 

5 

6 

7 
007 

8 

9 

0 

I  '00013 

• 

007 

003 

OOI 

000 

OOI 

003 

012 

019 

10 

027 

037 

048 

060 

073 

087 

103 

120 

138 

157 

20 

177 

199 

221 

244 

268 

294 

320 

347 

376 

405 

30 

435 

466 

497 

530 

563 

598 

633 

669 

706 

743 

40 

782 

821 

861 

901 

943 

985 

028 

072 

116 

^62 

5° 

i  '01207 

254 

301 

349 

398 

448 

498 

548 

600 

652 

60 

705 

758 

812 

867 

923 

979 

036 

093 

iS1 

210 

70 

i  '02270 

330 

390 

452 

5M 

576 

639 

7°3 

768 

833 

80 

899 

965 

032 

099 

768 

237 

306 

376 

447 

5^8 

90 

1-03590 

663 

736 

810 

884 

959 

035 

in 

188 

265 

100  11-04343 

422 

5oi 

*i  gram  of  ice  has  a  volume  of  i  .09008  cm. 3 


TABLE  VII. 

DENSITY  OF  MERCURY.* 
(At  the  Temperature  (t)  on  the  Normal  Hydrogen  Scale.) 


t° 

Density 

t° 

Density 

t° 

Density 

t° 

Density 

t° 

Density 

t° 

Density 

0 

I3-595° 

I 

•5925 

6 

.5802 

ii 

•5679 

16 

•5556 

21 

•5433 

26 

•5310 

2 

.5901 

7 

•5777 

12 

•5654 

J7 

•5531 

22 

.5408 

27 

.5286 

3 

.5876 

8 

•5753 

13 

.5629 

18 

•5507 

23 

•5384 

28 

.5261 

4 

•5851 

9 

•5728 

14 

•56°5 

!9 

•5482 

24 

•5359 

29 

•5237 

5 

13-5827 

10 

I3-5703 

15 

i3-558i 

20 

J3-5457 

25 

13-5335 

3° 

13.5212 

*Thiesen,  Guillaume,  International  Congress  of  Physics,  Paris,  1900. 
25 


386 


EXPERIMENTAL   CHEMISTRY. 


TABLE  VIII.     BAROMETRIC  CORRECTIONS. 

(a)  REDUCTION  or  BAROMETER  READINGS  TO  o°.     CORRECTIONS  FOR  THE  EXPAN- 
SION OF  THE  MERCURY  AND  THE  SCALE. 
The  following  corrections*  are  to  be  subtracted. 


Tem- 
pera- 
ture. 

BRASS  SCALE. 

GLASS  SCALE. 

OBSERVED  HEIGHT. 

OBSERVED  HEIGHT. 

c. 

740 

750 

760 

770 

780 

740 

750 

76o 

770 

780 

I 

2 
3 

4 

mm. 

•  12 

mm. 

.  12 
•25 

•37 
•49 

mm. 

.  12 

•25 
•37 
•SO 

mm. 
•13 
•25 
•38 
•50 

mm. 
•  13 

•5i 

mm. 

•38 
•Si 

mm. 
•13 
.26 
•39 
•52 

mm. 
•  13 
.26 
•39 
•53 

mm. 
•13 
•27 
.40 
•S3 

mm. 
•13 

•27 
•41 
•54 

5 

.  60 

.61 

.62 

•63 

.64 

.64 

.65     .66     .67 

.68 

6 

I 

9 

.72 
•8s 
•97 
1-09 

•73 
.86 
.98 

I-  10 

•74 
•87 
•99 

I  •  12 

•75 
.88 

I-OI 

I-I3 

.76 
.89 

1-02 

•77 
.90 

1-02 

.78 
•  9i 
1-04 
i-i7 

•79  ;    -80 
•92  !    -93 
I-  05  !   1-07 

I.  l8      1-20 

•  81 

•95 
i.  08 

I  •  21 

10 

I-  21 

I-  22 

1-24 

1.26 

I-  27 

1.28 

1.30    1-31    1-33 

i-35 

ii 

12 
13 
14 

1-33 
1-45 
1-57 
1-69 

••35 

i'47 
i'59 
1-71 

1-49 

1-61 
i-73 

1.38 
1^76 

1-40 
'•S3 
1-65 
1-78 

I-4I 

i-53 
1.66 
1.79 

1-43 
1-56 
1-69 
1.81 

i-45 
1-58 
1-71 

1-84 

1.46 
i.  60 
1-73 
1.86 

1.48 
1.62 
i-75 
1.89 

IS    1-81 

1-83 

1.86 

1-88 

1-91 

1.92 

1-94 

1-97      2-OO 

2  •  O2 

O  00<I  Ov 

1-93 
2-05 
2-  17 
2-29 

1-96 
2.08 

2-  2O 
2-32 

1-98 

2-  10 

2.23 
2-35 

2-01 

2.13 

2-  26 
2.38 

2.03 
2-16 

2-29 
2.41 

2-05 
2-17 
2-30 
2-43 

2-07 

2-  20 
2-33 
2-46 

2-10     2-13 
2-23      2.26 

2.36    2-39 
2-49    2.53 

2-16 

2-  29 

::0 

2O 

2.41 

2-44 

2-47 

2-SI 

2-54 

2-56 

2-59    2-62  i   2-66    2-69 

21 
22 
23 

24 

2-53 
2-65 
2-77 
2-89 

2.56 
2.69 

2.81 
2-93 

2-60 
2-72 
2-84 
2-97 

2.63 
2.76 

2-88 
3-oi 

2.67 
2-79 
2-92 
3  -05 

2-68 
2.81 
2.94 
3-06 

2-72 
2.85 
2.98 
3-  ii 

2-  76 

2.89 

3-02 

3-15 

2.79 
2.92 
3-o6 
3-19 

2.83 

2.96 
3-10 

3-23 

25 

3-oi 

3-05 

3-09 

3-13 

3-17 

3-19 

3-23 

3-28 

3-32 

3-36 

to  to  to  to 

VO  00<I  Ov 

3-  13 
3-25 
3-37 
3-49 

3-17 
3-29 
3-41 
3-54 

3-21 
3-34 
3-46 
3-58 

3-26 
3-38 
3-Si 
3-63 

3-33 
3-42 

III 

3-32 
3-45 
3-57 
3-70 

3-36 
3-49 
3-62 
3-75 

3-41 
3-54 
3-67 
3-8o 

3-45 
3-59 
3-72 
3-85 

3-50 
3-63 
3-77 
3-90 

30 

3-6i 

3-66 

3-71 

3-75 

3-8o 

3-83 

3-88 

3-93 

3-98 

4-03 

31 
32 

33 

34 

3-73 
3-85 
3-97 
4.09 

3-78 
3-90 
4.02 
4.14 

3-83 
3-95 
4-07 
4.20 

3-88 
4.00 
4-13 
4-25 

3-93 
4-05 
4-18 
4-31 

3-95 
4-08 

4-21 

4-33 

4-01 
4-14 
4-26 
4-39 

4.06 
4-19 
4-32 
4-45 

4-n 
4-25 
4-38 
4-51 

4.17 
4-30 
4-43 
4-57 

35 

4.21 

4-26 

4-32 

4-38 

4-43 

4.46 

4-52 

4-58 

4-65 

4.71 

*The  coefficient  of  expansion  of  mercury  is  0.0001813.  If  j3  is  the  coefficienct  of  expansion 
of  the  scale,  an  amount  equal  to  (o.oooi8i3-j3)ht  must  be  subtracted  from  the  observed  height 
(h)  to  obtain  the  reading  if  the  scale  and  mercury  were  reduced  to  o°  C.  (Glass,  ft  =  .000009; 
brass,  fi  =  . 00002;  steel,  )S  =  . 00012.) 

(b)  CORRECTION  FOR  VARIATION  IN  g. 


Latitude 

730 

740 

750 

760 

770 

780 

790 

mm. 

mm. 

mm. 

mm. 

mm. 

mm. 

mm. 

45° 

0 

o 

0 

o 

o 

0 

0 

40°  or  50° 
35°  or  55° 

•32 
•65 

III 

:i? 

•35 
.68 

111 

•35 
.70 

APPENDIX    II. 


387 


TABLE  IX. 

REDUCTION  TO  VACUUM  OF  WEIGHINGS  MADE  IN  AIR. 

If  a  body  of  density  (D)  has  an  apparent  weight  of  (G)  grams  when  weighed  in 
in  the  air,  its  weight  reduced  to  vacuum  is  G  +  Gk  grams,  k  is  computed  for  air  of 
density  .0012  and  for  brass  weights  of  density  8.4,  and  platinum  weights,  21.5. 

f  .0012         _    .0012    \ 

Go  =  G   I       ~D~     ^T~J"G(l+k) 


Correction  in 

Milligrams. 

Density  of  Body.     D. 

Brass  Weights. 

Platinum  Weights. 
d.  =  21  '5. 

k. 

k. 

0-50 

+  2-26 

+  2-34 

'55 

2  '04 

13 

•60 

•86 

'94 

65 

•70 

'79 

•70 

'57 

•66 

'75 

•46 

'55 

•80 

•36 

'44 

'85 

•27 

•36 

•90 

•28 

'95 

'  12 

'  21 

'0 

•06 

'14 

'  i 

'95 

'04 

'2 

•86 

'94 

'3 

•78 

•87 

•4 

"71 

•80 

'5 

•66 

'75 

•6 

•61 

•69 

'7 

;s6 

•65 

•8 

•62 

1-9 

'49 

•58 

2  '0 

•46 

'54 

2'5 

'34 

'43 

3'o 

•26 

'34 

3'5 

'20 

•29 

4'o 

•16 

•24 

5  ° 

'  10 

19 

6-0 

•06 

'14 

8-0 

+  '01 

•09 

IO'O 

'02 

•06 

15'° 

"06 

'03 

20  "o 

—  'OS 

•004 

EXPERIMENTAL   CHEMISTRY. 


TABLE  X. 

Volume  of  a  Glass  Vessel  at  Various  Temperatures. 
Correction  Factors  for  Calibrating  Glass  Vessels. 

In  calibrating  glass  vessels  by  weighing  them  filled  with  water  (or 
mercury)  at  a  known  t°,  it  is  necessary  to  correct  for  the  effect  of  tempera- 
ture on  the  density  of  the  liquid,  the  buoyant  effect  of  the  air,  and  the 
effect  of  temperature  on  the  volume  of  the  flask.  To  ascertain  the  true 
volume  (in  cm.3)  of  a  flask  filled  with  distilled  water  at  a  temperature  of 
20°  C.,  multiply  its  apparent  weight  (in  grams)  by  1.0028.  These  "fac- 
tors" assume  the  use  of  brass  weights  in  air  of  density  .0012  (approx.), 
and  that  the  coefficient  of  cubical  expansion  of  glass  is  .000025. 


/. 

VOLUME-  WATER. 

/. 

VOLUME-MERCUR  ¥. 

o 

ccm. 

o 

ccm. 

5 

i.ooi  49 

5 

0.073  647 

6 

.001  49 

6 

.073  658 

7 

.001  50 

7 

.073  669 

8 

•ooi  53 

8 

.073  681 

9 

.001  57 

9 

•073  693 

10 

i.ooi  63 

IO 

0.073  704 

ii 

.001  70 

ii 

.073  716 

12 

.001  78 

12 

.073  727 

13 

.001  88 

13 

•073  739 

14 

.001  99 

14 

•073  750 

15 

I.  O02  IO 

15 

0.073  762 

16 

.002  23 

16 

•073  773 

17 

.002  38 

17 

•073  785 

18 

.002  53 

18 

•073  797 

19 

.002  70 

19 

.073  808 

20 

I  .002  87 

20 

0.073  820 

21 

.003  06 

21 

•073  831 

22 

.003  26 

22 

•073  843 

23 

.003  47 

23 

•073  854 

24 

.003  69 

24 

.073  866 

25 

1.003  92 

25 

0.073  877 

26 

.004  15 

26 

.073  889 

27 

.004  39 

27 

.073  900 

28 

.004  65 

28 

.073  912 

29 

.004  92 

29 

•073  923 

30 

1.005  I9 

30 

0.073  935 

APPENDIX   II. 


.589 


TABLE  XI. 

VAPOR  TENSION  OF  WATER. 
Between  o°  C  and  35°  C. 


w 
w 
M 

o 

H 

Q 

VAPOR  TENSION  IN  MILLIMETRES  OF  MERCURY. 

TENTHS. 

.0 

•' 

.2 

•3 

•4 

•5 

-6 

-7 

.8 

•9 

0 

4-57 

4.60 

4.64           4-67 

4-  70 

4-74 

4-77 

4-80 

4-84        4-87 

i 

2 

3 

4 

4.91 
5-27 
5-66 
6-07 

4-94 
5-31 
5-  70 
6.  ii 

4.98 
5-35 
5-74 
6.15 

5-02 

5-39 
5-78 

6.20 

5-05 
5-42 
5-82 
6.24 

5-09 
5-46 
5-86 
6.28 

5-12 

5-50 
5-94 
6-33 

5-i6 
5-54 
5-94 
6-37 

5-20 

5-58 

5-99 
6-42 

5-23 
5-62 
6-03 
6.46 

5 
6 

2 

9 

6-Si         6.55 

6.60 

6.64 

6.69 

6-74 

6.78 

6-83 

6-88         6.92 

6.97 
7-47 
7-99 
8-55 

7.02 
7-52 
8.  os 
8.61 

7-07 
7-57 
8-io 
8.66 

7-  12 
7.62 
8.15 

8-72 

7-  i? 
7-67 

8.2! 
8.78 

7.22 
7-72 
8-27 
8.84 

7-26 
7-78 
8-32 
8.90 

7-3i 
7-83 
8-38 
8.96 

7-36         7-42 
7-88         7-94 
8-43         8-49 
9  •  02          9-08 

IO 

9*14 

9-  20 

9.26         9.32 

9-39 

9-45 

9-51 

9-58 

9-64          9.70 

ii 

12 
13 

14 

9-77 
10-43 
ii.  14 
11.88 

9-83 

io-  50 

II.  21 
IL96 

9.90 
10-57 
11.28 
12.04 

9.96 
10-64 
II-36 
12-12 

10-03 
io-  71 
H-43 
12.19 

10-09 
10-78 
11.50 
12-27 

io.  16 
10-85 
11.58 
12-35 

10.23 
io  •  92 
11.66 
12-43 

10-30 
10-99 
n-73 
12-51 

10-36 
11-07 
ii.  81 
12-59 

15 

12-67 

12-  76 

12.84 

12-92 

13-00 

13.09 

13-17 

13-25 

13-34 

13.42 

16 

ii 

19 

13-Si 
14.40 
15-33 
16.32 

13.60 
14-49 
15-43 
16.42 

13-68 
14-58 
15-52 
16.52 

13-77    ; 
14-67 
15-62 
16.63 

13-86 
14-76 
15-72 
i6.73 

13-95 
14-86 
15-82 
16-83 

14.04 
14-95 
15-92 
16.94 

14.12 

14.21 

14-30 
15-23 

l6-  22 
17.26 

16.  02 
17-04 

16.  12 
17-15 

20 

I7-36 

17-47 

I7-58 

17.69 

17-80 

17.91 

18-02 

18.13 

18.24 

18.35 

21 
22 
23 

24 

18-47 
19.63 
20-86 
22-15 

18-58 
19-75 
20-98 

22-  29 

18-69 
19-87 

21.  II 
22-42 

18.81 
19-99 
21.24 
22-55 

18-92 

20-  II 
2L37 
22.69 

19-04 

20-  24 
21  -50 
22.83 

19.  16 
20.36 
21-63 
22.96 

19.27 
20  •  48 

21.  76 
23.10 

19-39 
20.61 
21-89 
23-24 

19-51 

20-73 
22-02 
23-38 

25 

23-52 

23-66 

23-80     |     23.94 

24.08 

24-23 

24-37 

24-52 

24-66 

24-81 

26 

3 

29 

24.96 
26-47 
28-07 
29.74 

25.  io 

26.63 
28.23 
29.92 

25-25 
26-78 
28.39 
30.09 

25-40 
26-94 
28-56 
30.26 

25-55 
27-  io 
28.73 
30-44 

25-70 
27.  26 
28-89 
30.62 

25-86 
27.42 
29.06 
30-79 

26.01 
27-58 
29-23 
30-97 

26.16 
27.74 
29-40 
3I-I5 

26-32 
27.90 
29-57 
31-33 

30 

3i-5i 

31.69 

3I-87 

32.06 

32-24 

32-43 

32.61 

32.80 

32.99 

33-i8 

3i 

32 

33 

34 

33-37 
35-32 
37-37 
39-52 

33-56 
35-52 

37-58 

39-74 

33-75 
35-72 
37-79 
39-97 

33-94 
35-92 
38-00 
40-  19 

34-14 
36-13 
38-22 
40.41 

34-33 
36-33 
38-43 
40.64 

34-53 
36-54 
38-65 
40.86 

34-72 
36-74 
38.87 
41.09 

34-92 
36-95 
39-o8 
41-33 

35-12 
37-16 
39-30 
41-55 

35 

41-78 

42  •  02 

42-25 

42-48 

42-72 

42.96 

43-  19 

43-43 

43-6? 

43-92 

390 


EXPERIMENTAL    CHEMISTRY. 


TABLE  XII. 
VAPOR  TENSION  OF  WATER. 

FROM  o°C.  TO  ioo°C. 


Temp 

Tension 

Temp 

Tension 

Temp 

Tension 

Temp 

Tension 

o 

mm. 
4-569 

25 

mm. 
23-5!7 

5° 

mm. 
91.98 

75 

mm. 

288.76 

i 

4.909 

26 

24.956 

5i 

96.66 

76 

301.09 

2 

5.272 

27 

26.471 

52 

IOI-55 

77 

313-85 

3 

5-658 

28 

28.065 

53 

106.65 

78 

327-  °5 

4 

6.069 

29 

29.744 

54 

111.97 

79 

340-73 

5 

6.507 

30 

31-  51 

55 

117.52 

80 

354.87 

6 

6.972 

31 

33-37 

56 

123.29 

81 

369-51 

7 

7.466 

32 

35-32 

57 

129.31 

82 

384.64 

8 

7.991 

33 

37-37 

58 

I35-58 

83 

400.29 

9 

8.548 

34 

39-52 

59 

142.  10 

84 

416.47 

10 

9.140 

35 

41.78 

60 

148.88 

85 

433-  19 

ii 

9.767 

36 

44.16 

61 

155-95 

86 

450.47 

12 

10.432 

37 

46.65 

62 

163.29 

87 

468.32 

13 

11.137 

38 

49.26 

63 

170.02 

88 

486.76 

14 

11.884 

39 

52.00 

64 

178.86 

89 

505-81 

15 

12.674 

40 

54.87 

65 

187.  10 

90 

525-47 

16 

i3-510 

4i 

57-87 

66 

I95-67 

9i 

545-77 

17 

14-395 

42 

61.02 

67 

204.56 

92 

566.71 

18 

15-330 

43 

64-31 

68 

213.79 

93 

588.83 

J9 

16.319 

44 

67.76 

69 

223.37 

94 

610.64 

20 

17-363 

45 

71.36 

70 

233-3I 

95 

633.66 

21 

18.466 

46 

75-13 

71 

243-62 

96 

657.40 

22 

19.630 

47 

79.07 

72 

254-3° 

97 

681.88 

23 

20.858 

48 

83.19 

73 

265.38 

98 

707-13 

24 

22.  152 

49 

87.49 

74 

276.87 

99 

733.i6 

*Taken  from  Ostwald's  Manual  of  Physico-Chemical  Measurements. 


APPENDIX   II. 


391 


TABLE  XIII. 

VAPOR  PRESSURE  OF  MERCURY. 
(Ramsay  and  Young, — Jour,  of  the  Chem.  Soc.,  1886,  p.  37.) 


Temperature. 

Vapor  Pressure, 
mm. 

Temperature. 

Vapor  Pressure, 
mm. 

40° 

o'OoS 

1  80° 

S'535 

50 

•015 

190 

12-137 

60 

'029 

200 

17-015 

70 

•052 

210 

23-482 

80 

•092 

220 

31-957 

90 

•160 

230 

42-919 

100 

•270 

240 

56"9!9 

no 

'445 

250 

74-592 

120 

•719 

260 

96-661 

130 

i'i37 

270 

123-905 

I4O 

i"763 

280 

!57"378 

ISO 

2-684 

290 

198-982 

1  60 

4'oi3 

300 

246-704 

170 

5'9°4 

392 


EXPERIMENTAL  CHEMISTRY. 


TABLE  XIV. 

BOILING  TEMPERATURE  (0  OF  WATER  AT  BAROMETRIC  PRESSURE  (b). 
(After  Regnault.) 


b 

t 

b 

/ 

b 

t 

b 

t 

b 

t 

680 

96.92° 

700 

97.72° 

720 

98.49° 

740 

99.26° 

760 

100.00° 

68  1 

.96 

01 

•75 

21 

•53 

41 

.29 

61 

.04 

682 

97.00 

02 

•79 

22 

•57 

42 

•33 

62 

.07 

683 

.04 

°3 

.83 

23 

.61 

43 

•37 

63 

.  ii 

684 

.08 

04 

.87 

24 

•65 

44 

.41 

64 

•15 

685 

.12 

°5 

.91 

25 

.69 

45 

•44 

65 

.18 

686 

.16 

06 

•95 

26 

.72 

46 

•48 

66 

.22 

687    .20 

07 

97-99 

27 

.76 

47 

•52 

67 

.26 

688 

.24 

08 

98.03 

28 

.80 

48 

•56 

68 

.29 

689 

.28 

09 

.07 

29 

.84 

49 

•59 

69 

•33 

690 

•32 

710 

.  ii 

73° 

.88 

75° 

•63 

77° 

.36 

691 

.36 

ii 

•IS 

3i 

.92 

5i 

.67 

71 

.40 

692 

.40 

12 

.19 

32 

•95 

52 

.70 

72 

•44 

693 

•44 

13 

.22 

33 

98.99 

53 

•74 

73 

•47 

694 

.48 

14 

.26 

34 

99  -°3 

54 

•78 

74 

•51 

695 

•52 

15 

•30 

35 

.07 

55 

.82 

75 

•55 

696 

-56 

16 

•34 

36 

.  ii 

56 

•85 

76 

•58 

697 

.60 

17 

•38 

37 

.14 

57 

.89 

77 

.62 

698 

.64 

18 

.42 

38 

.18 

58 

•93 

78 

•65 

699 

.68 

19 

.46 

39 

.22 

59 

.96 

79 

.69 

700 

97.72 

720 

98.49 

740 

99.26 

760 

100  .  OO 

780 

100.72 

TABLE  XV. 

BOILING  TEMPERATURE"  (/)  OF  WATER  AT  A  PRESSURE  OF  (a)  ATMOSPHERES. 

(Regnault.) 


t 

a 

/ 

a 

100°  C. 

I  .000 

180°  C. 

9.929 

121 

2  .025 

189 

12.  125 

134 

3.008 

199 

15.062 

144 

4.000 

213 

19.997 

152 

4.971 

225 

25-I25 

159 

5.966 

239 

27-534 

I7I 

8.036 

APPENDIX   II. 


393 


«4H 

III                11 

$ 

'oS'o                                                     ^"o 

>nCJ 

If! 

111 

OMO      o      •      •    O      •      •      •    H    H      H    CM    co 

11? 

JO    O          OOO      -OOOOtoOO      •      •    O      •    O      •      •    O    O    O 
CO    fO         OOO      •    O    O    ^t"CO    t~>»  OO      •      *O      *O      •      'O^"O 

Ill 

S^o       JOE*    .g.coco^M^g    .    -^   .-   .    .OH? 

3  r«3      . 

1:5.1 

to 

O                                                               Tf                 •"•    NO    NO     tOCO 

H      .M^MCOMCOCO.CO^^            ;    ^     ;    ^     ;  J    M 

0 

o                                 o 

3       CJ      4-J 

1-S-s 

oVS^NOv8No''d'^'^     ^    

3§a 

,5                                ..^.^.Q        ON 

lit 

o 

•**•  O                                                                                                        (N    O    O    rO 
tOtotoO          HCNifO               votoco             NOIO        to         •^•ONfOON 

O    O    f^.  t^^  ro  fO  f^  f^OGO    O  CO    ON  O    ON  fONO    ON  M  CO    ^f  ^    O  NO    M    CO 
<N    ON  t^  t^OO    ON  <N  CO     OOO    <N  NO    ON  T|-  NO  CO    ro  M  CO    roJ5    rOCO    W  NO 

u 

°0 

o. 

^2 

OJ 

> 

ti 

:  :  gJU  :::::::::::  :^|  :::::: 

(_ 

s^ 

SQ^IU       '  TJ  T3              .CL,W..                     .. 

*• 

*Critica 

394 


EXPERIMENTAL   CHEMISTRY. 


TABLE  XVII. 

COMPOSITION  OF  THE  AIR  BY  VOLUME. 
(Average.) 


Vols.  per  1000. 

Nitrogen 769 . 5000 

Oxygen 206 . 5940 

Aqueous  vapor 14 .  oooo 

Argon  * 

Carbon  dioxide 

Hydrogen 

Ammonia 

Ozone 

Nitric  acid 


9.3700 
0.3360 
o. 1900 
o . 0080 
0.0015 
o . 0005 


IOOO.OOOO 


*The  other  four  elements  of  the  argon  group  constitute  about  .012  parts  in  a  1000. 
Small  quantities  of  solids  are  also  present,  e.  g.,  ammonium  nitrate,  ammonium  car- 
bonate, sodium  chloride,  dust,  etc. 


TABLE  XVIII. 

DIFFUSION  OF  GASES. 
(Graham,  1834  ) 


Density 

Square  root 

I 

Actual 
velocity  of 

Gas. 

(air-  1) 

of  density 

V  Density 

diffusion  by 
experiment 

Hydrogen 

o  '  06926 

o  '  2632 

3*  77Q4. 

V8* 

Marsh  gas 

O  '  <"XQ 

o  '  74.76 

•  9-?7r 

I  '  344. 

Steam  

o  '  62  3? 

o  '  7806 

'2664 

Carbon  monoxide 

0*0678 

0*0837 

'016$ 

I  '  1149 

Nitrogen    

O'Q7I3 

w  y"«>/ 
o  *o8c;6 

'014.7 

I  '0143 

Ethylene 

O  '  O?8 

0*9889 

'  OI  12 

I  '  0191 

Nitric  oxide 

•  O2Q 

i  *  1096 

o  '  9808 

Oxvgen    . 

•  ioc6 

I  '  O^I  1 

O  '  OCIO 

0*0487 

Sulphuretted  hydrogen   . 

•  1912 

I  '  OQ  1  4. 

o  '9162 

O'QC 

Nitrous  oxide  

•  C27 

I  '  23^7 

0*8092 

0*82 

Carbon  dioxide  

*  C2QOI 

I  '2^6i? 

o  '  8087 

0*812 

Sulphur  dioxide  

2  '247 

I  '  4QO  I 

0*667  I 

o'68 

APPENDIX   II. 


395 


TABLE  XIX. 
Specific  Heat— Atomic  Heat. 

Dulong-Petifs  Law.   A=    6'4' 


Elements. 

S* 

A 

SxA 

Lithium  Li 

O  04.1 

7 

6  6 

Beryllium     ...                              .  .Be 

o  408 

7      7 

Boron  (amorphous)  B             .    .  .    . 

O.2C.4 

II  . 

o  •  / 

2.8 

Graphite   \ 

0,174 

}» 

2.1 

Diamond  j 
Sodium  Na 

0,143 
O.27O 

J 

27 

i-7 

6.7 

Magnesum.  .                                    Me   . 

O.24.CC 

24.4 

">.Q 

Aluminum  Al        

O,2O2 

27  . 

tr  .  cr 

Silicon  (cryst.)  .  .                               Si 

o,i6c; 

28 

4.6 

Phosphorus.  .                                   P 

0,180 

7T  . 

^•0 

Sulphur  (rhombic)  S 

0,178 

72  . 

c.  .  7 

Potassium  K 

o  166 

TO 

6.< 

Calcium  Ca        

0,170 

4O. 

"    0 

6.8 

Chromium  .  .               Cr 

0,100 

C2  . 

C.  .  2 

Manganese                                       Mn 

O  122 

err 

6.7 

Iron  Fe             

O,II2 

^.Q 

6.7. 

Cobalt  Co 

O  IO7 

CO 

6.7. 

Nickel  .                                             Ni 

O  1  08 

c8    7 

6  4 

Copper  Cu                  .  . 

O,OQ7. 

ou  •  / 
67.6 

e.  .0 

Zinc  Zn 

O.OQ3 

6?    4. 

6  i 

Gallium                                            Ga 

O  O70 

7O 

r    r 

Germanium  Ge 

O,OC,7 

72.  C 

cr  .4 

Arsenic  (cryst.)  As 

O,o82 

7cr  . 

6.2 

Selenium  (cryst  )                              Se 

o  080 

7O   2 

6  4 

Bromine  (solid)  .   Br     .        .... 

0,084 

70.7 

6  7 

Zirconium  Zr 

0,066 

QO.  6 

6  o 

Molybdenum  Mo  

0,072 

06. 

6.9 

Ruthenium  Ru 

0,061 

101  .  7 

6.7. 

Rhodium  .  .                                     Rh 

o.o<c8 

103  . 

6.0 

Palladium  Pd  

o,oc,o 

106.  c 

6.3 

Silver  .  .        .       Ag              .... 

o,o<c6 

107  .0 

6.0 

Cadmium                                         Cd 

O.OC.4 

112  .4 

6.0 

Indium                                               In 

O  OC.7 

lie 

6.4 

Tin  .   Sn     

0,0^4 

IIQ. 

6.4 

Antimony.  .  .  .                                   Sb 

O.OC.2 

1  2O.  2 

6.2 

Tellurium                                         Te 

0.047 

127.6 

6.0 

Iodine                                                 I 

O  OC.4 

126.0 

6.8 

Lanthanum  ...    La  

O,O4C. 

138.0 

6.2 

*  Richter. 


396 


EXPERIMENTAL   CHEMISTRY. 

TABLE  XIX.— Continued. 


Elements. 

S* 

A 

SxA 

Cerium  Ce  

0,04^ 

I4O.  2 

6.2 

Didymium'j'  Di 

O.O4C 

142 

6  4 

Tungsten                                        \V 

o  011 

184 

6  i 

Osmium  Os  

^J^OO 

O,O3I 

101  . 

6.1 

Iridium                                             Ir 

O.O32 

103 

6  3 

Platinum                                          Pt 

O  O32 

IQ4    8 

6  3 

Gold       ...          Au          

O,O^2 

107  .  2 

"  •  o 

6.3 

Mercury  (solid)                               Hg 

O  O32 

2OO 

6  4 

Thallium                                          Tl 

O  O3  3 

2Q4    I 

6   7 

Lead              .                                   Pb 

w>^oo 
O,O3I 

206  o 

6  4 

Bismuth                                           Bi       . 

O  O3O 

208 

6    3 

Thorium  .           Th  

O,O27 

232  .  <C 

w  •  o 

6.4 

Uranium                                         Ur          .  . 

o  027 

238  q 

6  4 

*  Richter. 


fDidymium  =  Pr.  and  Nd. 


APPENDIX   II. 


397 


TABLE  XX. 
Heat  of  Formation.* 

(Arranged  from  Thomsen's  Thermo-chemische  Untersuchungen.) 
(a)     Chlorides,  Bromides  and  Iodides. 


Formula. 

Chlorides. 

Bromides. 

Iodides. 

(anhyd)f  (aq.)J 

(anhyd.)   (aq.) 

(anhyd.) 

(aq.) 

A1R3 
SbR3 
AsRs 
BaR3 
BiR3 
CaR2 
CdR2 
CoR2 
CuR 
CuR2 
AuR 
AuR3 
HR 
FeR2 
FeR3 
PbR2 
LiR 
MgR2 
MnR2 
Hg2R2 
HgR2 
NiR2 
KR 
AgR 
NaR 
SrR2 
TIR 
TIR3 
SnR2 
SnR4 
ZnR2 

160900 
91400 
71400 
194200 
90600 
183800 
93200 
76400 
32800 
51600 
5800 
22800 

22OOO 
82OOO 
96OOO 
82700 
93800 
I5IOOO 
III9OO 
652OO 

54500 
74OOO 
105600 
29300 
97600 
184500 
48500 

237700 

129600 

204900 

70300 

159300 

47100 
169400 

12600 

144000 

196300 

174400 

2OI2OO 
962OO 
948OO 

154900 
75200 

179400 
75600 
72900 

107600 
48900 

14900 
47900 
42500 

24900 
32500 

16200 

62700 

40800 

27200 
39300 
999OO 
127700 
75900 
IOI3OO 
186900 
I28OOO 

8800 
8400 

55°° 

28300 
78000 

—6300 

13100 
47600 

64400 
79900 

544oo 
91300 
165000 
106100 

39600 
61200 

76100 
134600 
75700 

50900 
4190 

31100 
25600 

51200 
93700 
IOI200 

71800 
90200 

41400 
75000 

953oo 
27700 
85700 
157700 
41400 

80100 
13800 
69100 

96500 
195700 
38400 
89OOO 
SlIOO 
I57OOO 
II280O 

85600 
173800 

70300 
143400 

30100 

56400 

10500 

80800 
I272OO 
97200 

759oo 

90900 

49200 

60500 

*  Smoothed  values  have  been  taken. 

f  Compound  in  usual  state  of  aggregation. 

|  Compound  formed  in  dilute  aqueous  solution. 


398 


EXPERIMENTAL    CHEMISTRY. 


(b)  Sulphates,  Nitrates  and  Carbonates. 


Formula. 

Sulphates. 

Formula. 

Nitrates. 

(anhyd.)|     (aq.) 

(anhyd.)  |     (aq.) 

A12(S04)3 
BaS04 
CaSO4 
CdS04 
CoS04 
CuSO4 
H2S04 

878900 

337500 
33240 

22IIOO 
233900 
182500 
192900 
24OOOO 
2l62OO 
334100 
3O22OO 
249800 

Ba(N03)2 
Ca(NO3)2 
Cd(N03)2 
Co(N03)2 
Cu(N03)2 
HNOo 
Fe7Fo3)2 
Pb(N03)2 
LiN03 
Mg(N03)2 
Mn(NO3)2 
HgN03 
Hg(N03)2 
Ni(N03)2 
KN03 
AgN03 
NaNO3 
Sr(N03)2 
T1NO3 
Zn(N03)2 

219800 
216700 

I2I2OO 
II9300 
92900 
41600 

2I52OO 
2207OO 
115800 
II4300 
82200 
49IOO 
II9500 
97900 
II22OO 
206300 
147500 
38900 
67100 
II32OO 
III400 
23300 
IO62OO 
2I52OO 
48lOO 
132300 

336800 
231800 
230400 
198300 
2IO7OO 
235600 

FeSO4.7H2O 
PbSO4 
Li2S04 
MgS04 
MnSO4 

NiS04 
K2SO4 
Ag2S04 
Na2SO4 
SrSO4 
T12S04 
ZnSO4 

105500 
III  6OO 
2I057O 

340100 
323000 
263500 

2293OO 
337200 
162700 
32900 

120760 
II94OO 
28700 
III2OO 
219800 
58lOO 
138200 

344500 
167200 
328500 
330800 
22IOOO 
23OOOO 

2I27OO 
248OOO 

BaCO3 
CdCO3 
MnCO3 
Ag2C03 
SrC03 

28l3OO 
l8l900 
2IO80O 
I2I300 
279600 

CaCO3 
PbC03 
K2C03 
Na2CO3 

269200 
I682OO 
279500 
27IOOO 

286OOO 
276500 

APPENDIX    II. 


399 


(c)  Oxides,  Hydroxides  and  Sulphides. 


Formula. 

Oxides. 

Formula. 

Hydroxides. 

(anhyd.) 

(aq.) 

(anhyd.)      (aq.) 

As2O3 
BaO 

CaO 
CuO 
Cu2O 
H2O 
(vapor) 
(liquid) 

20°   C. 

PbO 
Li20 
HgO 
Hg20 
K20 
Ag20 
Na20 
SrO 

154600 
130400 
14500 
37100 
40800 

57061 
68360 

50300 

147000 
158200 
163300 

Al(OH), 
Ba(OH)2 
Cd,0,H20 
Ca,02,H2 
Co,O,H2O 
Au203,3H20 
Fe,0,H20 
Fe2703,3H20 

LLO,H 
Mg,0,H20 
Mn,O,H2O 
Ni,O,H2O 
K,0,H 
NaOH 
Sr,02,H2 
Zn,O,H2O 

296900 
216300 
65700 

2266OO 

231700 

63400 
13200 
68200 
191100 

cal. 
cal. 

II74OO 

166500 

148900 
94700 
60800 
103900 

IO2OOO 
216400 
82600 

22OOO 
24860 
99100 
5900 
99800 
128400 

164500. 

116400 
I  I  l8oO 
226lOO 

155200 
157700 

A12S3 
BaS 
CaS 
CuS 
Cu2S 
H2S 
PbS 
Li2S 
MgS 
HgS 
K2S 
Ag2S 
Na2S 
SrS 

124400 
99500 
92OOO 
IOOOO 

18200 
2700 
18400 

BaS2H2 
CaS2H2 
CdS,nH20 
FeS,nH2O 
Co,S,nH2O 
Ni,S,nH20 
LiSH 
MgS2H2 
ZnS,nH2O 
KSH 

NaSH 
SrS2H2 

I24IOO 
125300 

107100 

II02OO 

32400 
2l8oO 
19700 
17400 

7300 

II32OO 
I  IOOOO? 

64100 
IIO8OO 

79600 

16800 

102400 

33000 

88400  ? 
99200? 

39600 

111300 

63100 

58500 
II9700 

55000 

101900 
104700 

400 


EXPERIMENTAL    CHEMISTRY. 


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402 


EXPERIMENTAL   CHEMISTRY. 


TABLE  XXIII. 
Composition  of  Some  of  the  Important  Alloys. 


Gold  coin  (U.S.,  Fr.,  Ger.) 

Gold  coin  (Great  Britain) 


/  Gold 900 

\  Copper I0o 

/  Gold 916.6 

\?0PPer 83.4 


Silver  coin  (U.  S.) 

Silver  coin  (Great  Britain) 

Silverware  (sterling) 

Bronze 

Gun-metal 

Bell-metal 

Speculum-metal 

Aluminum  bronze.. 


Managanese  bronze.  .  . 

Phosphorus  bronze  ................. 


Silicon  bronze 


Red  brass. . 
White  brass 


German  silver 


Silver 900 

Copper ioo 

Silver 925 

Copper 75 

Silver 925 

Copper 75 

Copper 93-5-95 

Tin 6-4 

Zinc 0.5-1 

Copper ioo 

Tin I0 

Copper 78 

Tin 22 

67 

r  33 

topper 90-95 

Aluminum 10-5 

Copper 90 

Managanese 10 

Copper 90 

Tin 9 

Phosphorus 5~-8 

Copper 90 

Tin 

Silicon 

Copper 

Zinc 

Copper 

Zinc 

Copper 

Zinc 

Nickel 

Tin.. 


Britannia-metal. . 


Hard  pewter 


9 

8 

90 

10 

65 

35 

5° 

25 

25 

ioo 

Antimony 8 

Bismuth i 

Copper 4 

Tin 92 

Lead..  8 


APPENDIX   II. 

TABLE  XXIII.— Continued. 


403 


Tin, 


Solder  (soft).. 
Queen's  metal, 

Type-metal 


so*  pewter (Lead.:::: 

Tin 

Lead 

Tin 

Antimony. 
Bismuth  . 

Lead 

Tin 

Lead 

Antimony. 

Fusible  metal  (" Rose's  Metal"),  ....      l~m '' '  ' 

fuses  at  94°  C £•       lu' ' 

I  Bismuth. . 

(Tin 

Fusible  metal  ("  Wood's  Alloy"),  .... 
fuses  at  65°  C 

Babbitt  metal,. 


82 

18 

(varies)  50 

(varies)  50 


Lead 

Bismuth. . 
Cadmium. 

Tin 

Lead 

Antimony. 
Copper. . . 


25 


50 

12 

50 
13 

40 

13 

I 


404  EXPERIMENTAL   CHEMISTRY. 

TABLE  XXIV. 
Scale  of  Hardness. 

Mineralogists  have  found  it  convenient  to  select  a  number  of  minerals 
for  the  comparison  of  hardness.  They  designate  the  hardness  on  a 
scale  of  10. 


No.  i.  Talc.     Scratched  easily  by  the  finger-nail. 

No.  2.  Gypsum.     Scratched  with  difficulty  by  the  nail. 

No.  3.  Calcite.     About  as  hard  as  a  copper  coin. 

No.  4.  Fluorite.     Slightly  harder  than  a  copper  coin. 

No.  5.  Apatite.     Scratched  easily  by  knife-point. 

No.  6.  Orthoclase.     Scratched  with  difficulty  by  knife-point. 

No.  7.  Quartz.  Is  not  scratched  by  knife-point. 

No.  8.  Topaz  or  beryl. 

No.  9.  Corundum.     Similar  to  "emery." 

No.  10.  Diamond. 


TABLE  XXV. 
Color  Scale  of  Temperature. 

This  table  is  the  result  of  an  effort  to  interpret  in  terms  of  thermometric 
readings,  the  common  expressions  used  in  chemistry  in  describing  tem- 
peratures. It  is  obvious  that  the  values  are  only  approximations. 


Color 

Temperature. 

Incipient  red  heat. 

5oo°-55o0 

Dark  red  heat. 

65o°-75o0 

Bright  red  heat.                                            85o°-95o° 

Yellowish-red  heat. 

1050°-!  150° 

Incipient  white  heat.                                  i25o°-i35o° 

White  heat.                                                  I45o°-i55o° 

APPENDIX    II.  405 

TABLE  XXVI. 
Indicators. 


Phenolphthalein —  Alcohol  solution — i  part  solid  in  100 

(A  weak  acid.)*  parts  of  60  per  cent,  alcohol.     Col- 

C2oH14O4?±  H-  +  C20H13O'4  orless  by  acids;  also  by  CO2;  red 

(Colorless)                   (red)  by  alkalies. 


Methylorange —  Water  solution — i  part  solid  to  1000 

(A  weak  base.)  parts  of  water.  Red  by  mineral 

/N  =  N  —  C6H4N(CH3)2  acids;  not  affected  by  CO2;  yellow 
C6H4 

\SO3Na  color  by  alkalies. 


Litmus — Water  solution.  Red  by  acids;  blue  by  .alkalies. 


Cochineal —  Alcoholic  solution, — 3  grm.  of  powder 

in  400  parts  H2O  and  100  parts  al- 
cohol. Yellowish-red  by  acids; 
violet  by  alkalies. 


*Stieglitz,  Jour.  Amer.  Chem.  Soc.  (Nov.  1903). 


TABLE  XXVII. 

Soap- Bubble  Solution  No.  I. 


Newth  suggests  the  following  formula  for  preparing  soap  solutions: 
"Ten  grams  of  sodium  oleate  and  400  cm.3  of  distilled  water  are 
placed  in  a  stoppered  bottle,  and  allowed  to  stand  until  the  oleate  has 
dissolved,  without  warming;  100  cm.3  of  pure  glycerin  are  then 
added,  and  the  mixture  after  being  well  shaken  is  allowed  to  stand 
in  the  dark  for  a  few  days.  The  clear  solution  is  then  carefully 
decanted  or  siphoned  into  a  clean,  stoppered  bottle,  one  drop  of 
strong  ammonia  solution  added.  If  kept  in  the  dark  and  not  ex- 
posed to  the  air,  this  solution  may  be  preserved  for  years." 


406  EXPERIMENTAL    CHEMISTRY. 

Soap-Bubble  Solution  No.  II. 

The  following  formula  will  be  found  to  give  excellent  results: 
Place  about  100  grm.  of  shavings  of  the  best  castile  soap  in  a  liter 
flask;  add  nearly  a  liter  of  distilled  water.  Shake  the  mixture  until 
a  saturated  solution  is  obtained,  then  allow  it  to  settle.  To  two 
volumes  of  the  clear  soap  solution  add  one  volume  of  glyerin. 
Keep  in  a  stoppered  bottle  in  the  dark. 


TABLE  XXVIII. 
Graduated  Solutions. 

(a)  Fehling's  Solution. 

This  solution  is  used  frequently  in  making  a  quantitative  determina- 
tion of  sugar.  As  it  decomposes  on  keeping,  it  is  best  kept  ("  pre- 
pared") in  the  form  of  two  separate  solutions:  (i)  34.639  grm. 
of  crystallized  copper  sulphate,  made  up  to  500  cm.3  with  water; 
(2)  173  grm.  of  Rochelle  salt  and  60  grm.  of  sodium  hydroxide  made 
up  to  500  cm.3  with  water.  For  use,  equal  volumes  of  the  two  solu- 
tions are  mixed.  10  cm.3  of  the  deep  blue  solution  thus  obtained  are 
completely  reduced  by  0.05  grm.  of  dextrose,  or  by  0.0475  grm-  °f 
sucrose,  after  inversion.  The  color  is  discharged  during  the  reduc- 
tion. 


(b)  Nessler's  Solution. 

This  solution  is  used  in  estimating  ammonia  (free).  "Dissolve  17 
grm.  of  mercuric  chloride  (pulverized)  in  300  cm.3  of  water,  and  35 
grm.  of  potassium  iodide  in  100  cm.3  of  water.  Pour  the  mercuric 
chloride  solution  into  the  potassium  iodide  until  a  permanent  red 
precipitate  is  formed.  Add  a  20  per  cent,  solution  of  sodium  hydrox- 
ide until  the  volume  of  the  mixed  solution  amounts  to  one  liter. 
Keep  this  solution  in  a  cool,  dark  place — portions  of  the  solution 
to  be  removed  as  needed.  It  is  necessary  to  "  sensitize  "  the  solution 
before  it  is  ready  for  use;  this  is  accomplished  by  adding  slowly  a 
saturated  solution  of  mercuric  chloride,  with  constant  stirring, 
until  a  permanent  red  precipitate  forms :  Allow  the  solution  to  stand 
until  the  solids  subside,  or  filter.  It  is  now  ready  for  use  and  should 
have  a  light,  straw-yellow  color.  The  solution  deteriorates  by 
standing." 


APPENDIX   II. 


407 


TABLE  XXIX. 

Percentage  and  Specific  Gravity  of  Solutions  at  15°  C. 
(a)  Sulphuric  Acid  (Lunge  and  Isler). 


Specific 
Gravity. 

Per  cent. 
H2  SO4. 

Specific 
Gravity. 

Per  cent. 
H2  S04. 

Specific 
Gravity. 

Per  cent. 
H2  SO4. 

1  .000 

o.oo 

1.320 

41.50 

1  .640 

71.99 

I  .010 

i-57 

1-33° 

42.66 

.650 

72.88 

1.020 

3-03 

1.340 

43-74 

.660 

73.64 

1  .030 

4-49 

•350 

44.82 

.670 

74-51 

1.040 

5-96 

.360 

45.88 

.680 

75-42 

1  .050 

7-37 

•37° 

46.94 

.690 

76.30 

1.  060 

8.77 

.380 

48.00 

.700 

77.17 

1  .070 

10.  19 

•390 

49.06 

1  .710 

78.01 

1  .080 

ii  .60 

.400 

50.11 

1  .720 

78.92 

1  .090 

12.99 

.410 

5I-I5 

1-730 

79.80 

I  .  TOO 

14-35 

.420 

52-15 

1.740 

80.68 

I  .  IIO 

i5-7i 

i-43° 

53-n 

1-750 

81-56 

1  .  1  20 

17  .01 

1.440 

54-07 

1.760 

82.44 

1.130 

18.31 

i-45o 

55-03 

1.770 

83-32 

1  .  140 

19.61 

i  .460 

55-97 

1.780 

84.50 

1.150 

20.91 

1.470 

56.90 

1.790 

85.70 

1  .  1  60 

22  .  19 

1.480 

57-83 

1  .800 

86.90 

1  .  170 

23-47 

1.490 

58.74 

1.810 

88.30 

1.180 

24.76 

1.500 

59-70 

1.820 

90.05 

1  .  190 

26.04 

i  .510 

60.65 

1.825 

91  .00 

I  .  200 

27.32 

1.520 

61.59 

1.830 

92.  10 

I  .  210 

28.58 

i-53o 

62.53 

1-834 

93-05 

I  .  220 

29.84 

1-540 

63-43 

1-837 

94.20 

1.230 

31.11 

1-550 

64.26 

1-839 

95-0° 

I  .  240 

32.28 

i  .560 

65.08 

i  .840 

95.60 

1.250 

33-40 

i-57o 

65-90 

1.841 

97.00 

I  .  260 

34-57 

1.580 

66.71 

1.8415 

98.20 

I  .  270 

35-7i 

1-590 

67-59 

1.841 

98.7° 

1.280 

36.87 

i  .600 

68.51 

i  .840 

99.20 

I  .290 

38-03 

i  .610 

69-43 

1.839 

99.70 

1.300 

39-19 

i  .620 

70.32 

I  .310 

40.35 

1.630 

71.16 

408 


EXPERIMENTAL    CHEMISTRY. 


TABLE  XXIX.— Continued. 
(b)  Hydrochloric  Acid.* 


Specific 
Gravity. 

Per  cent. 
HC1. 

Specific 
Gravity. 

Per  cent. 
HC1. 

Specific 
Gravity. 

Per  cent. 
HC1. 

I  .OOO 

0.16 

1.070 

14.17 

.  140 

27.66 

1  .005 

I-I5 

1-075 

15.  16 

•145 

28.61 

I.  010 

2.14 

1.080 

16.  15 

.150 

29-57 

1  .015 

3.12 

1.085 

17-13 

•155 

30.55 

1  .020 

4-13 

.090 

i8.ii 

.160 

3I-52 

1.025 

5-15 

•095 

19.06 

-165 

32-49 

1.030 

6.15 

.  IOO 

20.  01 

.170 

33-46 

1-035 

7-15 

.105 

20.97 

I-I75 

34-42 

1  .040 

8.16 

.no 

21  .92 

1.180 

35-39 

1.045 

9.  16 

1.115 

22.86 

1.185 

36-31 

1.050 

10.  17 

I  .120 

23.82 

I  .  190 

37-23 

1-055 

ii.  18 

I.I25 

24.78 

I-I95 

38.16 

I  .060 

12  .  IQ 

I.I30 

25-75 

I  .  I2O 

39-n 

1  .065 

13.19 

I-I35 

26.70 

*  Lunge  and  Marchlewski  in  Zeit.  f.  angew.  Chem.  1891,  133. 
(c)     Nitric  Acid.     (Lunge  and  Rey.) 


Specific 
Gravity. 

Per  cent. 
HNO3 

Specific 
Gravity. 

Per  cent. 
HNO3 

Specific 
Gravity. 

Per  cent. 
HNO3 

I  .OO 

o.oo 

1.18 

29.38 

1-36 

57-57 

I  .Ol 

1.90 

1.19 

30.88 

1-37 

59-3'9 

1  .02 

3-70 

1  .20 

32-36 

1.38 

61  .  27 

1.03 

5-50 

I  .21 

33-82 

1.39 

63-23 

1  .04 

7.26 

I  .22 

35-28 

1.40 

65-30 

1-05 

8.99 

1.23 

.36.78 

1  .41 

67.50 

1.  06 

10.68 

-24 

38.29 

1.42 

69.80 

1.07 

12.33 

•25 

39.82 

1-43 

72.17 

1.  08 

13-95 

.26 

41-34 

•44 

74-68 

1  .09 

15-53 

.27 

42.87 

•45 

77.28 

I  .  IO 

17.11 

.28 

44.41 

.46 

79.98 

I  .  II 

18.67 

1.29 

45-95 

•47 

82.90 

I  .12 

20.23 

1.30 

47-49 

.48 

86.05 

I-I3 

21.77 

I-3I 

49.07 

•49 

89.60 

I.I4 

23-31 

1.32 

50-71 

1.50 

94.09 

I-I5 

24.84 

i-33 

52-37 

i-5i 

98.  10 

1.16 

26.36 

i-34 

54-07 

1-52 

99.67 

1.17 

27.88 

i-35 

55-79 

APPENDIX   II. 


409 


TABLE  XXIX.— Continued. 
(d)  Ammonium  Hydroxide.* 


Specific 
Gravity. 

Per  cent. 
NH4. 

Specific 
Gravity. 

Per  cent. 
NH3. 

Specific 
Gravity. 

Per  cent. 
NH3. 

0.882 

34-95 

0.922 

21  .  12 

0.962 

9-35 

0.884 

34.10 

0.924 

20.49 

0.964 

8.84 

0.886 

33-25 

0.926 

19.87 

0.966 

8-33 

0.888 

32.50 

0.928 

19-25 

0.968 

7.82 

0.890 

31-75 

0.930 

18.64 

0.970 

7-31 

0.892 

3l.°5 

0.932 

18.03 

0.972 

6.80 

0.894 

30.37 

0-934 

17.42 

0.974 

6.30 

0.896 

29.69 

0.936 

16.82 

0.976 

5-8o 

0.898 

29.01 

0.938 

l6.22 

0.978 

5-30 

0.900 

28-33 

0.940 

I5-63 

0.780 

4.80 

0.902 

27.65 

0.942 

15.04 

0.982 

4-3° 

0.904 

26.98 

0.944 

14.46 

0.984 

3.80 

0.906 

26.31 

0.946 

13.88 

0.986 

3-30 

0.908 

25-65 

0.948 

I3-3I 

0.988 

2.80 

0.910 

24.99 

0.950 

12.74 

0.990 

2.31 

0.912 

24-33 

0.952 

12.  17 

0.992 

1.84 

0.914 

23.68 

0-954 

II  .60 

0-994 

i-37 

0.916 

23-03 

0.956 

11.03 

0.996 

0.91 

0.918 

22.39 

0.958 

10.47 

0.998 

0-45 

0.920 

21-75 

0.960 

9.91 

1  .000 

o.oo 

*  Lunge  and  Wiernick  in  Ziet.  f.  angew.  Chem.  1889,  183. 


4io 


EXPERIMENTAL   CHEMISTRY. 


TABLE  XXX. 

Proportion  by  Weight  of  Absolute  Alcohol. 
(Mendeleef.)* 


Specific  gravity 
atis°C. 

Per  cent, 
of  alcohol. 

Specific  gravity 
at  15°  C. 

Per  cent 
of 
alcohol. 

Specific 
gravity 
at  15°  C. 

Per  cent, 
of 
alcohol. 

0.9991 

o-5 

0.9501 

34 

0.8773 

68 

0.9981 

I 

0.9491 

35 

0.8750 

69 

0.9963 

2 

0-9473 

36 

0.8726 

70 

0.9945 

3 

0-9455 

37 

0.8702 

7i 

0.9928 

4 

0.9436 

38 

0.8678 

72 

0.9912 

5 

0.9417 

39 

0-8655 

73 

0.9896 

6 

0.9397 

40 

0.8631            74 

0.9881 

7 

0.9377 

4i 

0.8607           75 

0.9867 

8 

0-9357 

42 

0.8582           76 

0.9853 

9 

0.9336 

43 

0.8558 

77 

0.9839 

10 

0.9316 

44 

0.8534 

78 

0.9826 

ii 

0.9294 

45 

0.8510 

79 

0.9813 

12 

0.9273 

46 

0.8485 

80 

0.9801 

13 

0.9251 

47 

o  .  8460 

81 

0.9789 

14 

0.9230 

48 

0.8435 

82 

0.9777 

15 

0.9208 

49 

0.8410 

83 

0.9765 

16 

0.9186 

50 

0.8386 

84 

0-9753 

i7 

0.9164 

5i 

0.8360 

85 

0.9741 

18 

0.9142 

52 

0-8335 

86 

0.9728 

i9 

0.9119 

53 

0.8309 

87 

0.9716 

20 

0.9097 

54 

0.8283 

88 

0.9704 

21 

0.9074 

55 

0.8257 

89 

0.9691                      22 

0.9052 

56 

0.8230 

90 

0.9678                      23 

0.9029 

57 

0.8203 

9i 

0.9665                      24 

0.9097 

58 

0.8176 

92 

0.9651                      25 

0.8983 

59 

0.8149 

93 

0.9637 

26 

0.8960 

60 

0.8120 

94 

0.9623                      27 

0.8937 

61 

0.8092 

95 

0.9608                      28 

0.8914 

62 

0.8063 

96 

0-9593               29 

0.8890 

63 

0.8034 

97 

0-9577               30 

0.8867 

64 

o  .  8004 

98 

0.9561               31 

o  .  8844 

65 

0.7973 

99 

0-9544               32 

0.8820 

66 

0.7942 

IOO 

0.9527 

33 

0.8797 

67 

*Pogg.  Ann.  138,  p.  103. 


APPENDIX    II. 


411 


TABLE  XXXI. 

Proportion  by  Volume  of  Absolute  Alcohol. 
(Mendelee/.)* 


ioo  volumes  spirits. 

ioo  volumes  spirits. 

ioo  volumes  spirits. 

Specific  gravity 
at  15.5°  C. 

Contain 
volumes  of 
alcohol. 

Specific  gravity 
at  15.5°  C. 

Contain 
volumes 
ot 
alcohol 

Specific 
gravity 

15-5°  c. 

Contain 
volumes 
of 
alcohol. 

I  .  OOOO 

O 

0.9604 

34 

0.8950 

68 

0.9985 

I 

0.9591 

35 

0.8925 

69 

0.9970 

2 

0-9577 

36 

0.8901 

70 

0.9956 

3 

0.9563 

37 

0.8876 

7i 

0.9942 

4 

0.9548 

38 

0.8851 

72 

0.9928 

5 

0-9534 

39 

0.8825 

73 

0.9915 

6 

0.9518 

40 

0.8800 

74 

0.9902 

7 

0.9503 

4i 

0.8774 

75 

0.9889 

8 

0.9486 

42 

0.8747 

76 

0.9877 

9 

0.9470 

43 

0.8721 

77 

0.9866 

10 

0.9454 

44 

0.8694 

78 

0.9854 

ii 

0.9436 

45 

0.8667 

79 

0.9844 

12 

0.9419 

46 

0.8640 

80 

0.9832          13 

o  .  9400 

47 

0.8611 

81 

0.9822 

14 

0.9382 

48 

0.8583 

82 

0.9811 

15 

0.9364 

49 

0.8554 

83 

o  .  9$q  i 

16 

0.9344 

50 

0.8525 

84 

0.9790 

17 

0-9325 

5i 

0.8496 

85 

0.9781 

18 

o.9305 

52 

0.8466 

86 

0.9771              19 

0.9285 

53 

0.8435 

87 

0.9761 

20 

0.9265 

54 

o  .  8404 

88 

0-9751 

21 

0.9244 

55 

0.8372 

89 

0.9741 

22 

0.9222 

56 

0.8340 

90 

o.973i 

23 

0.9201 

57 

0.8306 

9i 

0.9720 

24 

0.9180 

58 

0.8272 

92 

0.9709 

25 

0.9158 

59 

0.8236 

93 

0.9699 

26 

0.9139 

60 

0.8199 

94 

0.9688 

27 

0.9113 

6! 

0.8161 

95 

0.9677 

28 

o  .  9090 

62 

0.8121 

96 

0.9667 

29 

0.9067 

63 

o  .  8080 

97 

0.9654 

30 

0.9045 

64 

0.8035 

98 

0.9642 

31 

0.9022 

65 

0.7989 

99 

0.9630 

32 

0.8997 

66 

0.7939 

IOO 

0.9617 

33 

0.8974 

67 

*Pogg.  Ann.,  138,  230. 


412 


EXPERIMENTAL   CHEMISTRY. 


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New  York.     $5.00. 
Seidell,  Atherton.     Solubilities  of  Inorganic  and  Organic  Substances.     D.  Van  Nos- 

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Thorpe.     Dictionary  of  Applied  Chemistry.     Longmans,  Green  &  Co.,  N.  Y.     3  vols. 

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Watts,  Henry.     Dictionary  of  Chemistry.     Revised  by  Morley  and  Muir.     Longmans, 

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Ladenburg,  Dr.  A.  Lectures  on  the  History  of  the  Development  of  Chemistry  Since 
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Meyer,  Ernest  von.  History  of  Chemistry.  Trans,  by  G.  McGowan.  The  Mac- 
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Muir,  M.  M.  Pattison.  A  History  of  Chemical  Theories  and  Laws.  Wiley  &  Sons, 
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Muir,  M.  M.  Pattison.  Heroes  of  Science,  Chemists.  Thomas  Nelson  &  Sons. 
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Ramsey,  W.     Gases  of  the  Atmosphere.     The  Macmillan  Co.,  N.  Y.     $2.00. 

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Venable,  F.  A.  Development  of  the  Periodic  Law.  Chemical  Publishing  Co.,  Easton, 
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Bloxam.     Inorganic   and    Organic    Chemistry.     P.    Blakiston's    Son  &   Co.,    Phila. 

(New  Ed.) 

Gooch  and  Walker.     Outlines  of  Inorganic  Chemistry.     The  Macmillan  Co. 
Holleman.     Inorganic  Chemistry.     Trans,  by  Cooper.     John  Wiley   &  Sons.,  N.  Y. 

$2.50. 

Jones,  H.  C.     Principles  of  Inorganic  Chemistry.     The  Macmillan  Co.,  N.  Y.     $4.00. 
Mendeleeff,   Dimitry   Ivanovitch.     Principles   of   Chemistry.     Longmans,    Green    & 

Co.,  N.  Y.     2  vols.     $10.00. 
Newth,  G.  S.     Text-book  of  Inorganic  Chemistry.     Longmans,  Green   &  Co.,  N.  Y. 

$1.75. 
Ostwald,  Wilhelm.     The  Principles  of  Inorganic  Chemistry.     The  Macmillan  Co., 

N.  Y.     $6.00. 

Remsen,  Ira.     Chemistry  (Adv.  Course).     Henry  Holt   &  Co.,  N.  Y.     $2.80. 
Richter,  Victor  von.     Inorganic  Chemistry.     P.  Blakiston's  Son  &  Co.,  Phila.     $1.75. 
Roscoe  and  Schorlemmer.     Treatise  on  Chemistry  (Revised,  1905;  i  vol.  ready).     The 

Macmillan  Co.,  N.  Y. 
Smith,  Alexander.     General  Inorganic  Chemistry.     The  Century  Pub.  Co.,  N.  Y. 

415 


416  REFERENCE    BOOKS. 

Organic  Chemistry. 

Gattermann,  Ludwig.     Practical  Methods  of  Organic  Chemistry.     Trans,  by  Shober. 

The  Macmillan  Co  ,  N.  Y.     $1.60. 
Holleman,  A.  F.     A  Text-book  of  Organic  Chemistry.     Trans,  by  Walker.     John 

Wiley  &Sons,  N.  Y.     $2.50. 
Lassar-Cohn.     Laboratory  Manual  of  Organic  Chemistry.     Trans,  by  Smith.     The 

Macmillan  Co.,  N.  Y.     $2.25. 

Mann,  Gustav.     Chemistry  of  the  Proteids.     The  Macmillan  Co.,  N.  Y. 
Noyes,  W.  A.     Organic  Chemistry.     D.  C.  Heath  &  Co.,  Boston. 
Orndorff,  W.  R.     Laboratory  Manual  in  Organic  Chemistry.     D.  C.  Heath   &  Co., 

Boston.     35  cents. 

Perkin  and  Kipping.    Organic  Chemistry.    J.  B.  Lippincott  Company,  Phila.     2  vols. 
Remsen,  Ira.     Organic  Chemistry.     D.  C.  Heath  &  Co.,  Boston.     $1.20. 
Richter,  Victor  von.     Organic  Chemistry.     Trans,  by  Smith.     P.  Blakiston's  Son   & 

Co.     2  vols.     $6.00. 

Theoretical  and  Physical  Chemistry. 

Arrhenius,  Svante.     Theories  of  Chemistry.     Longmans,  Green  &  Co.,  N.  Y. 

Ames. — Scientific  Memoirs.     American  Book  Co.,  Cin. 

Dobbin,  L.  and  W.  J.  Chemical  Theory  for  Beginners.  The  Macmillan  Co.,  New 
York.  70  cents. 

Freund,  I.  The  Study  of  Chemical  Composition.  The  Cambridge  Press,  Cam- 
bridge, Eng. 

Findlay,  Alexander.     Practical  Physical  Chemistry.     Longmans,  Green  &  Co.,  N.  Y. 

Jones,  Harry  C.  Elements  of  Physical  Chemistry.  Theory  of  Electrolytic  Dis- 
sociation. The  Macmillan  Co.,  N.  Y. 

Le  Blanc,  Max.  Electro-Chemistry.  Trans,  by  Whitney  and  Brown.  The  Mac- 
millan Co. 

Meyer,  Oskar  Emil.  The  Kinetic  Theory  of  Gases.  Trans,  by  Baynes.  Longmans, 
Green  &  Co.,  N.  Y. 

Meyer,  Lothar.  Outlines  of  Theoretical  Chemistry.  Trans,  by  Bedson  and  Williams. 
Longmans,  Green  &  Co.,  $2.50. 

Morgan,  J.  L.     Elements  of  Physical  Chemistry.     John  Wiley  &  Sons.,  N.  Y. 

Nernst,  W.  Theoretical  Chemistry.  Trans,  by  R.  A.  Lehfeldt.  The  Macmillan 
Co.,  N.  Y.  •  4v,i 

Remsen,  Ira.     Theoretical  Chemistry.     Lea  Brothers   &  Co.,  Phila.     $2.00. 

Talbot  and  Blanchard.  The  Electrolytic  Dissociation  Theory  with  some  of  its  Appli- 
cations. The  Macmillan  Co.,  N.  Y. 

Walker,  James.  Introduction  to  Physical  Chemistry.  The  Macmillan  Co.,  N.  Y. 
$3.00. 

Wurtz.     The  Atomic  Theory.     D.  Apple  ton  &  Co.,  N.  Y.     $1.50. 

Physiological  Chemistry. 

Arrhenius,  Svante.     Immunochemistry.     The  Macmillan  Co.,  N.  Y. 

Gamgee.     Chemistry  of  the  Animal  Body.     The  Macmillan  Co.,  N.   Y.     2  vols. 

$9.00. 
Hammarsten,   Olof.     A  Text-book  of  Physiological  Chemistry.   Trans,   by  Mandel. 

John  Wiley  &  Sons.,  N.  Y.     $4.00. 

Industrial  Chemistry. 

Groves  and  Thorp.     Chemical    Technology.     P.    Blakiston's    Son     &    Co.,    Phila. 

4  vols.     $16.00. 

Thorpe,  Frank  H.     Outlines  of  Industrial  Chemistry.     The  Macmillan  Co.     $3.50. 
Sadtler,  S.   P.     Industrial  Organic  Chemistry.     J.   B.   Lippincott  Company,  Phila. 

$5.00. 


REFERENCE    BOOKS.  417 

Analytical  Chemistry. 

Bottger,  Wilhelm.     Qualitative  Analysis.     Trans,  by  Smeaton.     P.  Blakiston's  Son 

&  Co,  Phila. 

Classen,  Alexander.     Quantitative  Analysis.     George  Wahr,  Ann  Arbor,  Mich. 
Coblentz,  Virgil.     Volumetric  Analysis.     P.  Blakiston's  Son  &  Co.,  Phila. 
Dennis,  L.  M.     Hempel's  Gas  Analysis.     The  Macmillan  Co,  N.  Y. 
Fresenius.     Quantitative    Chemical    Analysis.     Trans,    by    Cohn.     Wiley     &    Sons, 

N.  Y.     2  vols.     $12.50. 

Leffman  and  Beam.     Food  Analysis.     P.  Blakiston's  Son  &  Co,  Phila. 
Newmann,    Bernhard.     Electrolytic    Methods    of    Analysis.     Trans,    by    Kershaw. 

Whittaker   &  "Co,  N.  Y. 
Olsen,  J.  C.     Quantitative  Chemical  Analysis.     D.  Van  Nostrand  Company,  N.  Y. 

$4.00. 

Sulton,  Francis.     Volumetric  Analysis.     P.  Blakiston's  Son  &  Co,  Phila.  $5.00. 
Stillman,  Thomas  B.     Engineering  Chemistry.     Chemical  Publishing  Co,  Easton,  Pa. 
Treadwell,  F.  P.     Analytical  Chemistry.     Tran.  by  Hall.     John  Wiley  &  Sons,  N.  Y. 

2   Vols.       $7-00. 

Miscellaneous. 

Cornwall,  H.  B.     Manual  of  Blowpipe  Analysis.     D.  Van  Nostrand  Co,  N.  Y. 
Dana,  E.  S.     Minerals  and  How  to  Study  Them.     John  Wiley  &  Sons,  N.  Y.     $1.50. 
Fournier,  E.  E.     The  Electron  Theory.     Longmans,  Green   &  Co,  N.  Y.     $1.50. 
Hardin:  W.  L.     Rise  and  Development  of  Liquefaction  of  Gases.     The  Macmillan 

Co,  $1.50. 

McMillan,  W.  G.     Electro-Metallurgy.     Charles  Griffin  &  Co,  London. 
Rutherford,  E.     Radio-Activity.     University  Press,  Cambridge,  Eng. 
Shenstone,  W.  A.     Methods  of  Glass-Blowing.     Longmans,   Green    &  Co,   N.   Y. 

50  cents. 

Tables. 

Landolt   and   Bornstein.     Physikalisch-chemische  Tabellen.     G.  E.  Stechert    &  Co, 

N.  Y. 
Ostwald,    Wilhelm.     Physico-Chemical    Measurements.     Trans,    by    Walker.     The 

Macmillan  Co,  N.  Y. 
Traube,  Dr.  J.     Physico-Chemical    Methods.     Trans,    by    Hardin.     P.    Blakiston's 

Son   &  Co,  Phila 


INDEX. 


Acetamide,  248 
Acetone,  248 
Acetates,  259 

solubility  of,  260,  400 
Acetylene,  234 
Acid*,  acetic,  248,  259 

arsenic,  220 

arsenious,  220 

boric,  272 

bromic,  160 

carbolic,  256 

carbonic,  230 

chloric,  155 

chloroplatinic,  368 

chromic,  345,  346 

dithionic,  190 

formic,  248 

fumaric,  250 

hexathionic,  190 

hydrazoic,  199 

hydriodic,  162 

hydrobromic,  157 

hydrochloric,  107,  108,  409 

hydrocyanic,  237 

hydrofluoric,  152,  153 

hydrofluosilicic,  153,  243 

hypobromous,  159 

hypochlorous,  154 

hyposulphurous,  190 

iodic,  163,  164 

maleic,  250 

manganic,  352 

metaphosphoric,  217 

metastannic,  335 

molybdic,  349 

muriatic,  see  Hydrochloric  acid,  107, 
1 08,  409. 

nitric,  208,  209,  210,  408 
oxidation  by,  210 
test  for,  212 

nitrosyl  sulphuric,  186 

nitrous,  212 

test  for,  212 

oleic,  262 

orthophosphoric,  216 

oxalic,  260,  261 

*A11  acids  are  listed  under  "Acids." 


Acid,  palmitic,  262 

pentathionic,  190 

perchloric,  155 

perchromic,  346 

periodic,  165 

persulphuric,  190 

phosphoric,  216 

picric,  256 

prussic,  see  Hydrocyanic  acid,  237 

pyrophosphoric,  217 

pyrosulphuric,  190 

pyrosulphurous,  190 

selenic,  191 

silicic,  241 

stannic,  338 

stearic,  262 

sulphuric,  184 

chamber  process,  185 
contact  process,  188 
dissociation  of,  187 
heat  of  formation  of,  189 
heat  of  solution  of,  186 
properties  of,  186 

sulphurous,  183 

tartaric,  260 

thiocyanic,  238 

thiosulphuric,  189,  190 

trithionic,  190 

telluric,  191 

tungstic,  350 
Acidic  salts,  116,  117 
Acidimetry,  nS 

and  alkalimetry,  285 
Acids,  81,  107 

activity  ("strength")  of,  141,  142 

characteristic  properties,  no 

"condensed,"  272 

definitions  of,  in,  117,  141 

general  properties  of,  no 

ionization  of,  145 

monobasic,  dibasic,  etc.,  117 

nomenclature  of,  in 

organic,   in,  248,  259,  260 

oxascids  and  hydrascids,  in 
Acids  and  bases,  interaction  of,  114 
Acids,  bases  and  salts,  104 


419 


420 


INDEX. 


Affinity,  chemical,  36,  51 
Affinity-constant,  120 
Air,  see  Atmosphere. 
Albumin,  249 
Alcohol,  absolute,  253 

ethyl,  252,  253 

methyl,  248 

table  of,  420,  421 

tests  for,  254 
Alcohols,  248 
Aldehyde,  248 

preparation  of,  258 
Alkalies,  114 
Alkaline  earth  metals,  306 

detection  of,  314 
Alkali  metals,  280 

detection  of,  290 
Alkalimetry  and  acidimetry,  285 
Allotropy,  definition,  38 
Alloys,  276,  277 

table  of,  402,  403 
Alum,  preparation  of  common,  330 

iron,  359 
Alumina,  332 
Aluminates,  329 
Aluminum,  328,  329 

properties  and  compounds,  329-331 
Aluminum  family,  271 

comparative  table,  275 
Aluminum  salts,  analytical  reactions  of, 

.  33i 

Aluminothermy,  329 
Alums,  classification  of,  331 
Amalgams,  316 

ammonium,  289 
Amethyst,  242 
Amides,  248 
Amines,  248 
Ammonia,  199 

preparation  and  properties,  200 

volumetric  composition,  203 

weight  of  a  liter  of,  202 
Ammonium,  288 

compounds,  201,  288 

hydroxide,  table,  409 

salts,  reactions  of,  289 
Analysis     or     decomposition     (type     of 

chemical  change),  definition,  28 
Analysis,  qualitative,  of  anions,  371 

cations,  370 

Anhydride,  meaning  of  term,  in 
Anhydrous,  definition,  78 
Anion,  126 

Anions,  names  of,  135 
Anode,  126 
Antimony,  221 

properties  and  compounds,  222-223 

table,  225 

test  for,  222 


Apatite,  formula,  315 

Aqua  regia,  221 

Aqueous  (vapor)  tension,  84 

table  of,  389,  390 
Argentum,  see  Silver,  292,  299 
Argon,  197 
Arrhenius,  124,  127,  129 

theory  of  electrolytic    dissociation, 

130,  131 
Arsenates,  220 
Arsenic,  218 

properties  and  compounds,  221 

table,  225 

tests  for  218,  225 
Arsenites,  220 

Arsenuretted  hydrogen  (arsine)  218 
Arsine,  218 
Asbestos,  314 

platinized,  68 
Atmosphere,  193 

composition,  etc,  193-198 

table  of,  394 
Atom,  definitions  of,  50 
Atomic,  heats,  table  of,  395 

hypothesis,  50 

volume,  174,  175 

weight,  50 

weights,  412 
Aurates,  304 
Avogadro's  rule,  49,  128 
Azurite,  305 

B 

Balanced  action,  122 
Balance,  chemical,  379 
Bancroft,  W.  D.,  quoted,  37 
Barite  (heavy  spar),  215 
Barium,  313 

properties  and  compounds,  313-314 

tests  for,  314 

Barometer,  Table  of  corrections  for,  386 
Baryta  water,  231 
Bases,  81,  113,  117,  142 

activity  (strength)  of,  142 

characteristic  properties,  113 

classification  of,  113-114 

general  properties,  113 

ionization,  table  of,  145 

common  ion,  from,  142 

monacid,  diacid,  117 

nomenclature,  114 

"strong,"  113 
Basic  salts,  116,  117 
Basicity,  117 
Basic  oxides,  114 
Bayley's  table,  173 
Bauxite,  332 
Becquerel  rays,  334,  350 
Beilby,  G.  T.,  133 


INDEX. 


421 


Bell-metal,  see  Table  of  Alloys,  402,  403 

Bending  glass,  instructions,  375 

Benzene,  246 

Berlin  (Prussian)  blue,  360 

Berthollet,  on  law  of  definite  proportions, 

45 

Berthollet's  laws  of  chemical  reaction,  122 
Berthelot's  proposed  third  law  of  thermo- 
chemistry, 39 

Berzelins'  "dualistic  theory,"  125 
Bessemer  process,  see  Iron,  356 
Bibliography,  415-417 
Bicarbonate  of  soda,  230 
Binary  compounds,  1 1 1 
Bismuth,  223 

properties  and  compounds,  223-224 

table  of  properties,  225 
Blast  lamp,  374 
Bleaching,  by  chlorine,  105 

by  sulphur  dioxide,   182 
Bleaching  powder,  154,  310 
Blue-stone  (blue  vitriol)  see  Copper  sul- 
phate 

Blow-pipe,  use  of,  240 
Blow-pipe  flame,  oxidation  with,  240 

reduction  with,  240 
Body,  term  defined,  18 
Boiling-point,  elevation  of,  98 

constants,  99 
Boracite,  272 

Borates,  solubility  of,  275,  400 
Borax,  272 
Borax  bead,  273 

colors  of,  274 
Boron,  271 

properties  and  compounds,  271-274 

table  of  properties,  275 

test  for,  274 
Boyle's  law,  20 
Brass,  see  Alloys,  276,  277 
Brimstone,  see  Sulphur,  176,  177 
Britannia  Metal,  see  Alloys,  276,  277 
Bromides,  solubility  of,  159,  400 
Bromine,  156 

properties  and  compounds,  156-160 

table  of  properties,  165 
Bronze,  335 

see    Table  of  Alloys,  402,  403 
Brookite,  333 
Bunsen  burner,  study  of,  338,  373 

flame,  239 

temperature  of  flame,  240 
Burettes,  376-377 
Butter,  262 


Cadmium,  316 

properties  and  compounds,  321-322 
salts,  analytical  reactions  of,  322 


Caesium,  290 
Calamine,  316 
Calcite,  315 
Calcium,  309 

properties  and  compounds,  309-312 

tests  for,  311 

Calculations  in  chemistry,  52 
Calomel,  323 
Calories,  kinds  of,  52 
Camphor  gum,  228 
Cane-sugar,  265,  268 

structural  formula  for,  251 
Caramel,  268 
Carbohydrates,  248 
Carbon,  227 

amorphous,  227 

disulphide,  236 

properties  of,  227-229 

halides  of,  236 

hydrides  of,  233-235 

oxides  of,  229-232 
Carbon  and  nitrogen,  236,  237 
Carbon  compounds,  some  common,  246- 

270 
Carbon  family,  226 

table  of,  244 
Carbonates,  230 

solubility  of,  231 

test  for,  230 
Carborundum,  243 
Carnallite,  308 
Cassiterite,  335 

Catalysis  (catalytic  action),  58,  69 
Catalyzer,  definition  of,  58 

action  of,  58,  59 
Cathode,  126 
Cation,  126 
Cations,  names  of,  135 
Caustic,  lunar,  303 
Caustic  alkalies,  114 
Cavendish,  54 
Cellulose,  265,  269 
Cement,  315 

hydraulic,  329 
Cerite",  326 
Cerium,  326,  333 
Chalcopyrite,  356 
Chalk,  315 

Chalybeate  water,  355 
Chamber  process,  185 
Charcoal,  227,  228 
Charles'  law,  20 

Chemical  action,  illustrations   of,  22,  26, 
27,  28,  29 

means  of  initiating  by  heat,  26,    32, 
36 

by  solution,  29 

Chemical  actions,  reversible,  119 
Chemical  affinity,  36,  51 


422 


INDEX. 


Chemical  changes,  varieties  of,  28,  29,  31, 

33 

characteristics  of,  26-30 
Chemical  energy,  36,  51 
Chemical  equilbrium,  119,  121 

laws  of,  122 

Chemical  equivalent,  48 
Chemical  phenomena,  22 
Chemical  properties,  33,  54 
Chemical  reactions,  complete,  119 

incomplete,  119,  121 

reversible,  119,  124 

laws  of,  122 
Chemistry  and  physics,  relation  between, 

17,  18 

Chemistry,  organic,  227 
Chemistry,  science  of,  17 
Chemistry,  thermo-,  38,  39 
China  ware  (porcelain),  327 
Chlorates,  1-55 

solubility  of,  155 

test  for,  155 
Chlorides,  164 

solubility,  154 

test  for,  347 
Chlorine,  104,  153 

properties  and  compounds,  104-110, 

I53~I56 

table  of  properties,  165 
Chloroform,  261,  262 
Chromates,  345,  346 

oxidizing  power  of,  346 
Chrome-alum,  347 
Chrome-yellow,  341 
Chromic,  anhydride,  346 
Chromites,  345 
Chromium,  344 

"passive",  344 

properties  and  compounds,  344-349 

salts,  analytical  reactions  of,  348 
Chromous  compounds,  344 
Chromyl  chloride,  347 
Cinnabar,  316,  323 
Clay,  329 
Coal,  distillation  of,  235 

gas,  235 

tar>  335 
Cobalt,  355,  363 

properties  and  compounds,  363-364 

salts,  reactions  of,  364 
Cobalt  amines,  363 
Cobaltite,  363 
Cobaltous    compounds,  363 
Coke,  235 

Colloidal  solution,  83,  242 
Columbium  (niobium),  343 
Combination  (synthesis),  29 
Combining  proportions,  laws  of,  45,  46 

volumes,  law  of,  48 


Combining  weights,  law  of,  48 
Combustion  of  carbon  in  oxygen,  heat  of, 

58 
Common    ion,    effect    of    adding,    ionic 

equilibrium,  146 
Complex  ions,  see  Ions 

salts,  144,  298 

Component,  definition  of  physical,  32 
Compound,  32 

endothermic,  39 

exothermic,  39 
Compounds,  binary,  in 

ternary,  in 
Concentration,  90 

expression  of,  90 
Concentration  (mass)  effect,  120 
Concentration  (ionic)  effect,  320 
Conductivity,  electrical,  135-138 

equivalent,  138 

interpretation  of,  131,  136,  137 
Conservation  of  energy,  21 

of  mass,  1 8,  43 
Constant,  affinity,  120 

electro-chemical,  134 

ion-product,  146 

ionization,  139 
Constants,  molecular  depression,  100 

molecular  elevation,  99 

physical,  table  of,  393 
Constituent,  definition  of  chemical,  32 
Constitution  of  substances,  249 

of  alcohol,  proof  of,  250 
Copper  (cuprum),  292 

analytical  reactions  of  salts  of,  297 

equivalent  of,  296 

properties  and  compounds  of,  292- 

299 

Copperas,  359 
Copper  group,  292 
Corks,  boring  of,  375 

rubber,  treatment  of,  376 
Corpuscles,  see  Electrons,  132 
Corrosive  sublimate,  323 
Corundum,  329 
Cotton,  269 

Critical  phenomena,  49 
Critical  solution  temperature,  95 
Cry  of  tin,  335 
Cryolite,  152 
Cryp to-crystalline,  241 
Crystallization,  water  of,  78,  79 

see  Hydration 

Cupric  compounds,  see  Copper 
Cuprous  compounds,  see  Copper 
Curie,  M.  and  Mme.,  on  radioactive  sub 

stances,  350 
Cyanates,  237 
Cyanides,  237 
Cyanogen,  236 


INDEX. 


423 


D 

Dalton's  law  of  partial  pressures,  62 

Davy,  Sir  Humphry,  112,  125 

Decomposition  (analysis),  28 
heat  of,  39 

Decrepitation,  78 

Definite  proportions,  law  of,  45 

Deflagration,  287 

Degrees  of  ionization,  table  of,  145 

Dehydration,  78 

Deliquescence,  78 

Density,  19 

tables  of,  384,  385 

Depression  constant,  molecular,  100 

Depression  of  the  freezing-point,  97,  98, 
100 

"Developer"  (in  photography),  301 

Dextrose,  270 

Dialysis,  242 

Diamond,  227 

Dichromates,  345,  346 

Didymium,  326 

Dielectric  constants,  132 

Diffusion  of  gases,  law  of,  67 
table  of,  394 

Diffusion  phenomena  in  solution,  83 

Dilution  formula,  Ostwald's,  138,  139 

Displacement  of  ions,  by  a  free  metal,  146 
by  a  free  non-metal,  147 
power  of  metals,  order  of,  147 

Dissociation  constant,  139,  142 

Dissociation,  electrolytic,  124-151 

hydrolytic,  see  Hydrolysis,  77,  80, 144 
in  solution,  proofs  of,  124-151 
of  compounds  by  heat  ,  123 

Distillation,  77 

Dobereiner's  triads,  168 

Double  decomposition  (metathesis),  33 

Double  (compound)  salts,  144,  298 

Dulong  and  Petit's  law,  52,  395 

Dyeing,  see  Mordants,  331 

Dynamic  condition  of  chemical  equilib- 
rium, 123 

Dynamite,  254 

Dysprosium,  412 


"Earth"  metals,  306 
Earths,  alkaline,  306 
Efflorescence,  78 
Eka -aluminum,    (gallium),    327 
Eka-boron  (scandium),  327 
Electrical  energy  and  chemical  energy,  148 
Electro-chemical  equivalent,  134 
Electrolysis,  definition  of,  126 
Electrolysis  and  electrical  equivalents,  134 
Electrolyte,  126,  135 

Electrolytic   determination   of   copper   in 
cupric  sulphate,  297 


Electrolytic  dissociation  theory,  130 
origin  and  development  of,  124 

Electromotive  force  (potential),  149 
series,  150 

E.  M.  F.  as  a  measure  of  the  tendency 
toward  chemical  reaction,  149,  150 

Electronic  theory,  40,  132 

Electrons,  132 

Electro-plating,  definition  of,  364 

Element,  definition  of,  37 

Element  versus  "elementary  substance," 

37 

Elementary  (simple)  substances,  31 
Elements,  acid-forming,  118,  152 

base-forming,  118,  276 

classification  of,  167 

grouping  of,  279 

metalic,  31,  118,  152 

non-metallic,  31,  118,  276 

table  of,  412 
Emery,  329 
Enantiotrophy,  324 
Endothermic  compounds,  39 

preparation  of,  301 

Energetics  of  chemistry,  note  on  the,  34 
Energy,  20 

bound,  37 

chemical,  36 

conservation  of,  21 

definition  of,  20 

form,  24 

free  or  available,  37 

internal,  21,  35 

forms  of,  21,  35 

transformation  of,  21 
"Energy  content,"  34 
Energy  and  chemical  changes,  29,  36 
Energy  and  matter,  relation  of,  22,  23 
English  measures,  382 
Equations,  kinds  of,  50 
Equilibrium,  119 

chemical,  119,  121,  122 

ionic,  146 

study  of  ionic,  320 
Equivalents,  definition  of,  48 

chemical,  48 

electro-chemical,  134 
Erbium,  326 

Esters  (ethereal  salts),  248,  254 
Ethers,  248,  257 

compound,  258 
Ethyl  acetate,  248,  254,  262 
Ethyl  alcohol,  248 

properties  and  preparation,  252,  253 

table   of  proportion   by   weight   and 
volume  of,  410,  411 

tests  for,  254 
Ethylene,  233,  246 
Ethylene  bromide,  233 


424 


INDEX. 


Europium,  412 
Exothermic  compound,  39 
Experiment,  term  denned,  40 
Extracts,  fluid,  253 


Fact,  term  denned,  40 
"Factors,"  term  denned,  33 
Faraday,  126 

"Faraday"    (96,580    coulombs    of    elec- 
tricity), 134 
Faraday's  laws,  126 
Fat,  262 
Fatty  acids,  254 
Fehling's  solution,  406 
Feldspar,  241 
Fermentation,  252 
Ferrates,  355 
Ferric  alum,  359 
Ferric  compounds,  see  Iron 
Ferricyanides,  237,  360 
Ferrocyanides,  237,  360 
Ferro-manganese,  351 
Ferrous  compounds,  see  Iron,  355 
Fertilizers,  311 

"Fixing,"  in  photography,  301 
Flame,  study  of  a,  238 

Bunsen,  238 

oxidizing  and  reducing,  239 
"Flashing-point,"  251 
Flash-light  powder,  307 
Flax,  269 

Flowers  of  sulphur,  177 
Flint,  241 

Fluorides,  solubility  of,  153,  400 
Fluorine,  152 

preparation  and  properties  of,  152-153 

table  of  properties  of,  165 
Fluorite  (fluor-spar)  152 
Fluor-spar,  152 
Flux,  277 
Fools'  gold,  356 
Formaldehyde,  259 
Formalin,  259 

Formation,  heat  of,  39,  397,  398,  399 
Formulae,  kinds  of,  51 

dualistic,  51 

empirical,  50,  249 

graphic  or  structural,  52,  250 

rational,  51,  250 

space,  52 
Franklinite,  356 

Freezing-point,  depression  of,  100 
Fructose,  268 
Fusion  of  ice,  heat  of,  79 


Gadolinite,  326 
Gadolinium,  326 


Galena,  336 

Gallium,  326,  327,  328 

Garnet,  245 

Gas,  coal,  234,  235 

laughing,  see  Nitrous  oxide,  205,  207 

illuminating,  235 
Gas  law,  49 
Gases,  critical  constants  of,  393 

diffusion  velocities  of,  394 

heat  of  solution  of,  93 
solubility  of,  96,  97 
Gasoline,  251 

Gay-Lussac's  law,  see  Charles'  law,  20,  49 
Gay-Lussac's  law  of  gaseous  volumes,  48 
Gelatin,  formula  of,  238 
Gelatinoids,  249 
Germanium,  334 

table  of  properties  of,  337 
German  silver,  364,  402 
Glass,  etching  of,  153 
Glass,  uranium,  350 

Glass  vessels,  correction  factors  for  cali- 
brating, 388 
Glass  working,  3 74-3 75 
Glauber's  salt,  290 
Glucinum,  306 
Glucose,  265,  267,  268 
Glucosides,  249 
Glycerides,  255,  262 
Glycerin,  255 
Gold,  304 

properties  and  compounds,  304-305 
Graham  and  Bunsen's  law  of  effusion,  67 

table,  394 

Goldschmidt's  process,  328,  329,  330 
Gram-molecule  (mole),  a  chemical  unit 

of  weight,  52 

Gram-molecular  volume,  52 
Grape-sugar,  265 
Graphite,  227 
Grotthus,  quoted,  125 
Green  fire,  314 
Greenockite,  316 
Guldberg  and  Waage,  law  of  mass  action, 

120 

Gun-cotton,  266 
Gun-metal,  composition  of,  402 
Gunpowder,  286 
Gypsum,  310 

H 

Halides,  152,  248 

Halite,  see  Sodium  chloride 

Halogen  family,  152 

Halogens,  table  of  relations   of,  164,  165 

Hardness  (water),  permanent,  312 

temporary,  231,  310 
Hardness,  scale  of,  404 
Heat  of  combustion  of  carbon,  58,  230 


INDEX. 


425 


Heat,  of  combustion,  58 

of  decomposition,  39 

of  formation,  39,  207 

tables  of,  397-399 

of  fusion  of  ice,  79 

of  ionization,  151 

of  neutralization,  143 

of  reaction,  39,  207 

of  solution,  92,  93 

of  vaporization  of  water,  79 
"Heat  sum,"  law  of  the   "constancy  of 

the"  (law  of  Hess),  39 
Heavy-spar,  315 
Helium  family,  197 
Hematite,  356 
Hemp,  269 
Henry's  law,  96 

Hess,    law    of    (second    law  of  thermo- 
chemistry), 39 

Heterogeneous  systems,  18,  25 
Hittorf,  quoted,  127 
Homogeneous  systems,  18,  25 
Homologous  series,  246 
Hiibnerite,  349 
Hulett,  quoted,  84 
Hydrates,  80 
Hydraulic  cement,  329 
Hydrazine,  199 
Hydrides,  ex.,  73,  247 
Hydrion,  properties  of,  141 

velocity  (absolute),  135 
Hydrocarbons,  246 

derivatives  of,  247,  248 
Hydrogele,  242 
Hydrogen,  64 

arsenuretted,  s»e  Arsine,  218 

ionic,  141 

nascent  (active),  70 

preparation  and  properties,  64-73 
Hydrogen  bromide,  see  Acid,  hydrobromic, 

J57 
Hydrogen  chloride,  108,  109 

see  Acid,  hydrochloric,  107,  409 
Hydrogen  dioxide  (peroxide),  70,  71 
Hydrogen  fluoride,  152,  153 
Hydrogen  halides,  154 
Hydrogen  iodide,  162 

dissociation  of,  123 
Hydrogen  peroxide,  70,  71 
Hydrogen  sulphate,  186 

see  Acid,  sulphuric 
Hydrogen  sulphide,  178,  179,  180 

see  Acid,  hydrosulphuric. 
Hydrolysis,  77,  80,  81,  144 

of  salts,  77,  144 
Hydrosole,  242 

Hydroxides,  definition  of,  113 
Hydroxide-ion    (hydroxidion),    properties 
of,  144 


Hydroxide-ion,    velocity    (migration)    of, 

J35 

Hydro  xyl,  113 
Hygroscopic,  196 
"Hypo,"  302 

Hypochlorites,  oxidizing  power  of,  154 
Hypothesis,  atomic,  40 

Avogadro's  (Rule),  49,  128 

definition  of,  41 

ionic,  130 

kinetic-molecular,  49 

Prout's,  167 

I 
Ice,  see  Water 

heat  of  fusion  of,  79 
Iceland  spar,  see  Calcite,  315 
Identity,  definition  of,  18 
Illuminating-gas,  235 
Indicators,  360 

table  of,  405 
Indigo,  154 
Indium,  328 
Induction  in  chemical  reaction,  period  of, 

164 

Inductive  capacity,  specific,   132 
Ink,  359 

sympathetic,  364 
Internal  energy,  21,  35 
Internal  rearrangement,  a  kind  of  chemi- 
cal change,  31 

Iodide  (tri-odide)  of  nitrogen,  301,  302 
Iodide,  starch,  161 
Iodine,  152,  160 

preparation  and  properties,  160-164 

table  of  properties  of,  165 

test  for,  161 

tincture  of,  161 
lodoform,  261 
Ion,  definition  of,  126,  130 

effect  of  adding  a  common,  146 

product,  146 
Ionic,  concentration,  146 

equilibrium,  146,  320 
Ionic  hypothesis,  124,  130 

development  of,  124-130 

objections  to,  124 

support  for,  134-151 
Ionic  substances,  131,  135 
Ionization,  130,  131,  136 

change  in  volume  during,  143 

constant,  139 

factors  which  influence,  136-140 

heat  of,  151 

per  cent,  or  degrees  of,  table  of,  145 

repression  of,  363 

scheme  of,  298 
lonogens,  131 

classification  of,  131 


426 


INDEX. 


Inogens,  charges  on,  number  of,  130, 
Ions,  chemical  conduct  of,  141 

color  of,  140 

complex,  144 

displacement  of,  146,  147 

nomenclature  of,  135 

color  of,  140 

source  of  electrical  charges  of,  131- 

J34 

speed  of  migration  of,  135 
Indium,  367 
Iron  (ferrum),  355 

alum,  359 

carbide,  356 

cast,  356 

compounds,  355-362 

electrolytic,  356 

kinds  of,  356 

ores,  356 

passivity  of,  357 

Pig,  356 

properties  of,  356-358 

pure,  356 

pyrites,  356 

salts,    analytical    reactions    of,    360, 
361 

spiegel,  356 

wrought,  356 
Iron  elements,  355 
Isatin,  formula,  of  154 
Isomerism,  249 
Isomers,  249 


Jasper,  241 
Jones,  H.  C.,  82 


K 


Kalion,  migration  velocity,  135 
Kaolin,  245,  329 
Kerosene,  fire  test  of,  251 
flashing-point  of,  251 
Ketones,  248 
Kindling  temperature,  56 
Kinetic-molecular  hypothesis,  49 
Kohlrausch,  127 
Krypton,  see  Helium  family,  197 


Lamp-black,  228 

Lanthanum,  326 

Laughing-gas,  see  Nitrous  oxide,  205,  207 

Laplace,  quoted,  39 

Lavoisier,  quoted,  39 

Law,  Berthollet's,  of  chemical  reaction,  122 

Berthelot's,    proposed    for    thermo- 
chemistry, 39 

Boyle's,  49 

Charles',  49 


Law,  Dalton's,  62 

definition  of,  40 

Dulong  and  Petit's,  395 

Faraday's  laws,  126 

gas,  49 

Gay-Lussac's,  48 

Graham's,  67 

Henry's,  96 

Hess's,  39 

Le  Chatelier's  (theorem),  37 

mass,  Guldberg,  and  Waage's,  120 

of  combining  weights,  48 

of  concentration  effect,  see  Mass  law 

of  conservation  of  energy,  21 

of  constancy  of  composition,  45 

of  definite  proportions,  45 

of  depression  of  freezing-point,   100, 
101 

of  diffusion,  67 

of  dilution,  138 

of  distribution,  96 

of  energetics,  first,  21 

of  gaseous  volumes,  48 

of  Hess,  39 

of  ion-product,  139 

of  matter,  18 

of  multiple  proportions,  46 

of  partial  pressures,  62 

of  partition,  see  Law  of  distribution 
96 

of  reciprocal  proportions,  48 

of  thermochemistry,  39 

of  vapor  pressure  of  solvents,  129 

Ostwald's  dilution,  138 

periodic,  166 

Raoult's,  100,  101,  129 

Van't  Hoff's,  87 
Lead,  333,  334,  336 

properties  and  compounds,  336-341 

salts,  analytical  reactions  of,  341 

sugar  of,  340 

table  of  properties,  337 

white,  341 
Lead-tree,  339 
Le  Blanc's  theory  of  the  source  of  the 

electric  charges  on  the  ions,  131 
Levulose,  (fructose),  268 
Light,  in  chemical  action,  29 

see  Photography,  301 
Lime,  309 
Lime-water,  309 
Limonite,  356 
Litharge,  340 
Lithium,  281 

Litmus,  as  an  indicator,  405 
Lodestone,  see  Magnetite 
Lothar  Meyer's  table,  170 
Luminosity,    of    Bunsen     burner    flame, 
cause  of,  239 


INDEX. 


427 


Lunar  caustic,  303 

Lyes,  330,  see  Alkalies,  114 

M 

Magnesia  alba  (hydrated  basic  corbonate 

of  magnesium),  308 
"Magnesia  mixture,"  217 
Magnesium,  306,  307 

properties  and  compounds    of,  307- 

3°9 

salts,  analytical  reactions  of,  308 
Magnesium  nitride,  196,  212 
Magnetite,  356 
Malachite,  305 
Manganates,  351,  352,  353 
Manganese,  351 

properties  and  compounds  of,  351- 

354 

Manganese  alloys,  351 
Manganese  bronze,  402 
Manganese  spar,  351 
Manganic  compounds,  351 
Manganin,  351 
Manganites,  351 
Manganous  compounds,  351 
Marble,  309 
Marsh-gas,  233 
Marsh's  test,  215,  216 
Mass,  action,  120,  121,  122,  222 

conservation  of,  18,  43 

effect,  120 
Matrix,  277 
Matter,  17,  18 
Mayer,  Julius  Robert,  21 
Meerschaum,  314 
Mendeleef's  table,  169 
Meniscus,  377 

Mercuric  oxide,  heating  of,  28 
Mercuric  chloramide,  324 
Mercurous  chloramide,  324 
Mercury,  316,  317 

densities  of,  table  of,  385 

properties  and  compounds,  322—324 

reactions  of  salts  of,  324,  325 

table  of  properties  of,  317 
Metal,  definition  of  a,  118 
Metallic  elements,  see  Metals 
Metals,  (base-forming  elements),  276 

"alkali,"  280 

"alkaline  earth,"  306 

chemical  properties  of  the,  277 

classification  of  the,  278 

earth,  326 

grouping  for  purposes  of  analysis  of 
the,  370 

occurrence  in  nature  of  the,  277,  278 

physical  properties  of  the,  276 
Metallurgy,  277 
Metamers,  249 


Metathesis,  33 
Methane  (marsh  gas),  233 
Methods  of  science,  40,  41 
Methyl,  247 
Methyl,  orange,  405 
Mica,  245 

Microcosmic  salt,  217 
Migration  of  ions,  135 

speed  of,  135 

Milk  of  lime,    see  Lime-water,  309 
Mineral,  definition  of,  277 
Minium,  340 
Mixture,  31,  32 
Mohr's  salt,  359 
Moissan,  152 
Molar,  solutions,  90 

volume,  52 

weight,  52 
Molecular,  equations,  50 

formulae,  50,  51 

hypothesis,  49 

weight,  50 

methods  of  determining  the,  52 
Molecule,  definitions  of,  49 
Molecules,  compound,  50 

monatomic,  51 

simple,  50 
Molybdenum,  349 

properties  and  compounds  of,  349 
Monazite,  333 
Mordants,  331 
Morse  and  Frazer,  87,  88 
Mortar,  310 
Multiple  proportions,  law  of,  46 

N 

Naphtha  (naptha),  251 
Naphthalene,  256 
Nascent  condition,  70 
Nascent  hydrogen,  70 
Natrion,  135 

Natrium,  see  Sodium,  in,  280 
"Negative,"  in  photography,  301 
Neodymium,  326 
Neon,  see  Helium  family,  197 
Nernst,  93,  132,  148 
Nessler's  reagent,  406 
Neutralization,  114,  117,  143 

complete,  117 

experimental  study  of,-  114 

heat  of,  117,  143 

partial,  117 

theory  of,  143 

thermochemistry  of,  143 

volume  change  of,  143 
New  elements  predicted  by  the  periodic 

system,  327 

Newland's  octaves,  168 
Niccolite,  364 


428 


INDEX. 


Nickel,  335,  364 

properties  and  compounds,  364-365 

reactions  of  salts  of,  365 
Nickelion,  365 
Nickel-steel,  365 
Nilson  and   Cleve   discovered   eka-boron 

(scandium),  327 
Niobium  (columbium),  343 
Nitrates,  211 

effect  of  heat  on,  211 

oxidizing  power  of,  211 

solubility  of,  213,  400,  401 

tests  for,  212 
Nitric,  anhydride,  206 

oxide,  205 

Nitride  of  magnesium,  196,  214 
Nitrites,  212 

solubility  of,  213,  400 

tests  for,  213 
Nitrogen,  192,  199 

compounds,  199-213 

preparation  and  properties,  192 

table  of  properties,  225 

thermochemistry  of  oxides  of,    205- 

208 

Nitrogen  family,  199 
Nitrogen  tri-iodide,  201 
Nitroglycerin,  254 
Nitrosyl  chloride,  211 
Nitrosyl-sulphuric  acid,  186 
Nitrous,  anhydride,  206 

oxide,  205,  207 
Nomenclature,  49-52 

of  acids,  bases  and  salts,  in,  114,  117 
118 

of  ionic  theory,  126,  130,  134-151 
Non-electrolytes,  126,  135 
Non-metal,  definition  of,  118 
Non-metallic     elements,      grouping     for 

purposes  of  analysis,  371 
Non-metals,  grouping  of,  167 
Normal,  salts,  117 

solutions,  90 
Noyes,  A.  A.,  quoted,  20,  85 

O 

Occlusion,  68,  69 

Octaves,  law  of,  168 

Oil,  illuminating,  see  Kerosene,  251 

Oil  of  vitriol,  see  Acid,  sulphuric,  184 

Oils,  262 

drying,  263 

essential  or  volatile,  262 

fatty  or  fixed,  262 

non-drying,  263 

semi-drying,  263 
Olein,  262 
Oleomargarine,  263 
Olivine,  314 


Onyx,  242 

Opal,  242 

Open-hearth  process,  356 

Ores,  277 

Organic  chemistry,  227 

Orthite,  326 

Orthoclase,  245 

Osmium,  367 

Osmosis,  85 

Osmotic  pressure,  85,  88 

definition  of,  85 
Osmotic  pressure  and  molecular  weights, 

87 
Osmotic  pressure  of  cane  sugar,  86 

effect  of  concentration,  86 

effect  of  temperature,  86 
Ostwald,  quoted,  59,  69,  119 
Oxalates,  solubility  of,  261,  400 
Oxidation,  54,  56 

defined  in  terms  of  the  ionic  hypoth- 
esis, 211 
Oxides,  54 

basic,  114 

classification  of,  57 
Oxidizing  flame,  240 
Oxygen,  54 

nascent,  107 

preparation  and  properties  of,  54-59 
Oxygen  and  hydrogen,  heat  energy  pro- 
duced when  they  combine,  73 
Oxygen  and  ozone,  different  amounts  of 

energy  in,  59 
Oxygen  family,  176,  191 
Ozone,  59,  176 


Palladium,  367 
Palmitin,  262,  269 
Paper,  269 

parchment,  265 
Paraffin,  251 
Paris,  plaster  of,  310 
Partial  pressures,  law  of,  62 
"Passive"  condition  of  elementary  sub- 
stances, 344,  357 
Perchlorates,  155 
Periodic  arrangement,  171,  172,  173 

of  Lothar  Meyer,  170 

of  Mendeleef,  169 
Periodic  law,  166 

Periods  of  induction  in  chemical  action,  1 64 
Permanganates,  351,  352 

oxidizing  power  of,  353 
Petit  and  Dulong's  law,  395 
Petroleum,  251 
Pewter,  339,  402,  403 
Pfeffer,  128 

Pfeffer's  measurement  of  osmotic  pressure 
86 


INDEX. 


429 


Phase,  definition  of  term,  24,  25 
Phase  rule,  37 
Phenol,  256 
Phenolphthalei'n,  257 
Phenolphthalein,  as  an  indicator,  405 
Phenols,  248 
Phenomenon,  18 

surface,  29,  83,  84 
Phosphates,  216,  217 

solubility  of,  218,  400 

tests  for,  218 
Phosphine,  214 
Phosphorus,  199,  213 

-bronze,  402 

chemical  relations,  table  of,  225 

properties  and  compounds  of,  213- 
218 

properties  of  red  and  yellow,  38 

see  Phosphates,  216,  217 
Photography,  301-302 
Physical  constants,  table  of,  393 
Physical  and  chemical  changes,  26 
Pig,  iron,  356 
Pipettes,  378 
Pitchblende,  350 
Plaster  of  Paris,  310 
"Platinized"  asbestos,  68 
Platinum,  366,  367 

catalytic  action  of,  69,  368 

properties  and   compounds  of,  367- 
369 

reactions  of  salts  of,  369 

"spongy,"  catalytic  action  of,  68 
"Platinum  black,"  367,  369 
Plumbago,  227 
Plumbites,  334 

Plumbum,  see  Lead,  333,  336 
Polymerization  of  water,  80 
Polymers,  249 
Polysulphides,  181 
Portland  cement,  315 
"Positive" -photography,  301 
Potash,  caustic,  114 

yellow  prussiate  of,  237 
Potassium,  280,  286 

properties  and  compounds,  286-288 

radical  present  in   the  salt  or  com- 
pound sought  for,  288 

salts  of,  see  under  the  acid 

solubility  of  compounds  of,  288 

tests  for,  288 
Potential,  149 

Potentials,  table  of  single,  150 
Powder,  smokeless,  266 
Praseodymium,  326 
Precipitates,  terms  used  in  describing  the 

structure  of,  282 

Precipitation,    ionic,    theory   and   formu- 
lation of,  146 


Prediction,  definition  of,  41 
Prediction  of  new  elements,  327 
Pressure,  critical,  49 

osmotic,  84,  85,  87,  88 

partial,  62 

solution,  84,  147 

vapor,  62,  84 
Priestly,  54 
Product,  concentration,  146 

ion,  146 

solubility,  146 
"Products,"  chemical,  33 
Propane,  246 
Properties,  body,  19 

chemical,  18,  54 

definition  of,  18 

general,  18,  19 

physical,  18 

reaction,  54 

specific,  18,  19 

Proust   established   law   of  definite   pro- 
portions, 45 
Prout's  hypothesis,  167 
Prussian  blue,  360 
Pseudo-solutions,  83,  242 
Purification,  of  sodium  carbonate,  283 

of  sodium  chloride,  283 
Purple  of  Cassius,  305 
Pyrites,  356 
Pyrolusite,  350 


Qualitative  analysis,  cations,  371 

cations,  370 

Quantitative  relationships,  43 
Quantivalence,  see  Valence 
Quartz.  241 
Quartz-glass,  241 
Quicklime,  309 
Quicksilver  (mercury),  316 

R 

Radicals,  acid,  in 

organic,  247 
Radio-activity,  350 
Radium,  307,  314 
Ramsay,  193,  197 
Raoult's  laws,  100,  101,  129 
Rare  earth  metals,  326 
Reaction,  33 

heat  of,  39 
Reactions,  balanced,  122 

complete,  119,  121,  122 

incomplete,  119,  122 

reversible,  119 
Realgar,  218 
Red  fire,  312 

Red  heat,  temperature  corresponding  to, 
404 


430 


INDEX. 


Red  lead,  340 

Red  phosphorus,  38,  213 

Reducing  flame,  240 

Reduction,  in  terms  of  ionic  hypothesis, 

211 

Reversible  reactions,  119,  121,  122,  183 

Rhodium,  366 

Rinmann's  green,  320 

Richards,  T.  W.,  20,  31,  39,  40 

Rochelle  salt,  260 

Rock  crystal,  see  Quartz,  241 

Rock  salt  (halite),  see  Sodium  chloride 

Rose's  fusible  metal,  403 

Rubidium,  290 

Ruby,  329 

Rule,  Avogadro's,  49,  128 

definition  of,  40 
Rust,  306 

Rusting  of  iron,  356 
Ruthenium,  366 
Rutile,  333 


Saccharose,  265,  268 
Sal  ammoniac,  289 
Sal  soda,  291 
Salt,  common,  283 

definition  of,  117,  143 

electrolysis  of  a,  114 

Glauber's,  290 

Rochelle,  260 

impurities  of  common,  283 
"Salting  out"  of  soap,  263 
Salts,  114 

acid,  116,  117 

anhydrous,  78 

basic,  1 1 6,  117 

classification  of,  116,  117 

complex,  144,  298 

compound,  144,  298 

dehydrated,  78 

double  or  compound,  144,  298 

essential,  262 

ethereal,  248,  262 

fixed,  262 

hydrolysis  of,  144 

ionization  of,  per  cent,  of,  145 

nomenclature  of,  118 

normal,  116,  117 

solubilities  of,  400,  401 
Samarium,  326 
Sand,  153,  241 
Saponification,  264 
Sapphire,  329 
Saturated  solution,  83,  84,  85,  90 

making  of  a,  83 

theory  of,  83,  84 

Saturation,     a    term     used    in    organic 
chemistry,  247 


Scandium,  326 

Scheele,  54 

Science,  definition  of,  41 

development  of  a,  40,  41 

methods  of,  41 

natural,  41 
Sciences,  abstract,  42 

biological,  41,  42 

classification  of,  41,  42 

concrete,  42 

descriptive  or  special,  42 

general  or  speculative,  42 

physical,  41 

Scientific  versus  systematic  method,  41 
Selenite,  310 
Selenium,  190 
Series,  electromotive,  150 

in  organic  chemistry,  246 

in  periodic  system,  169,  171 

solution  tension,  147 
Serpentine,  245 
Shot,  making  of,  336 
Shot-metal,  336 
Siderite,  356 

Siemens-Martin  process,  356 
Silica,  153,  240 
Silicates,  242,  243 

solubility  of,  243 

test  for,  243 
Silicon,  226,  240 

compounds  of,  242,  243 

preparation  and  properties  of,  241 

table  of  properties,  244 
Silver  (argentum),  292,  299 

argenticyanide,  303 

compounds,  of  299-304 

electrolytic  deposition  of,  299 

preparation   from   an  alloy  of  pure, 
300 

properties  of,  299 

reactions    of    salts     of,    analytical, 

3°4 

sterling,  300,  402 

"tree,"  299 
Slag,  277 

Slaked  lime,  309,  310 
Smaltite,  363 
Smithsonite,  316 
Soap,  definition  of  a,  264 

hard,  263,  264 

preparation  of,  263 

properties  of,  264 

salting  out  of,  263 

soda,  263 

soft,  264 

Soap-bubble  solutions,  405,  406 
Soda,  ash,  282 
Soda,  caustic,  114 

washing,  291 


INDEX. 


431 


Sodium  (natrium),  in,  280,  281 

aluminate,  330 

cobaltic  nitrite,  285 

hydroxide,  113,  282 

properties  of,  112,  113,  281 

solubility  of  salts  of,  285 

salts  and  other  compounds,  282-286 

see  also  under  the  acid  radical 
present  in  the  salt  or  compound 
sought  for. 

zincate,  319 
Solder,  soft,  335 
Solubilities,  tables  of,  400,  401 
Solubility,  data,  400,  401 

effect  of  pressure  on,  93,  96 

effect  of  solvent  on,  96 

effect  of  surface  on,  83,  84 

effect  of  temperature  on,  88,  89,  95.  97 

explanation  of  degrees  of  solubility, 
82 

expression  of,  90 

measurement  of,  89 

mutual,  94 

of  gases,  96,  97 

of  liquids,  94 

of  pairs  of  liquids,  94 

of  solids,  82,  88,  89,  90,  91,  93 

product,  146 
Solute,  definition  of,  82 

distribution   of  solute   between   two 

immiscible  solvents,  95 
Solution,  characteristics  of  a,  82 

colloidal,  83,  221,  242 

definition  of  a,  83 

heat  of,  92,  93 

in  two  solvents,  95,  96 

modern  theory  of,  124 

partial  explanation  of  the  mechanism 
of,  83,  84 

process  of,  chemical  or  physical,  102 

temperature,  critical,  95 

tension,  84,  146,  147,  148,  149 

tension  of  metals,  146,  147 
Solutions,  82 

boiling-points  of,  98 

concentrated,  82,  90 

dilute,  82 

freezing-points  of,  98-101 

ionic,  equilibrium  of,  146 

kinds  of,  82 

lowering  of  the  vapor  tension  of,  97 

molar,  90 

normal,  90 

pseudo-(colloidal),  83,  242 

saturated,  82,  88 

standard,  90 

supersaturated,  89 
"  terminology  of,  90 

vapor  tension  of,  97 


Solution  tension,  electrolytic,  149 

hypothesis  of,  84 

lowering  of  the,  101 

of  the  metals,  147 

series,  150 

Solvent,  definition  of,  82 
Solvents,  immiscible,  103 
Soot,  228 
Space  formulae,  52 
Specific  gravities,  19,  173,  175 
Specific  heats  of  elements,  395 
Specific  inductive  capacity  of  water,  132 

of  various  solvents,  132 
Specific  volume,  19 
Spectroscope,  figures  of,  280,  281 
Speed  of  reaction  affected  by,  catalyzers, 
58,  68,  69,  1 88 

concentration,  120,  146,  222,  320 

solution,  29,  141 

surface  effects,  29,  66,  69,  307 

temperature,  26,  28,  32,  33,  123,  188 
Sphalerite,  316 
Spiegel  iron,  356 
Spirits,  253 

"Sponge,"  platinum,  68 
Stannates,  334 
Stannic  chloride,  338 
Stannic,  compounds,  see  Tin,  333,  336 
Stannum,  see  Tin,  333,  336 
Starch,  266 

test  for,  266 

theory  of  formation  of,  269 
Starch  iodide,  161 
Steam,  see  Water 
Stearin,  262 
Steel,  356 

nickel,  365 
Stibine,  222 
Strontianite,  315 
Strontium,  306,  312 

analytical  reactions  of  salts  of,  312, 

313 

compounds,  312-313 

ion,  313 

properties  of,  312 

Structural  or  graphic  formulae,  52,  250 
Structure,  molecular,  249,  250 

see  Constitution  of  substances,  249 
Sublimate,  corrosive,  323 

see  Mercuric  chloride 
Sublimation  of  iodine,  160 

of  salts  of  mercury,  323 
Substance,  compound,  definition  of,  32 

definition  of,  18 

simple  or  elementary,  definition  of, 

3*>  37 
Substitution,  236 

products,  236 
Sugar,  barley,  268 


432 


INDEX. 


Sugar,  beet-,  268 

cane-,  265,  268 

grape-,  265 

inversion  of  cane-,  265,  268 

Fehling's  test  for,  265 

invert,  268 

see  Saccharose 

structural  formula  for  cane-,  251 
Sugars,  classification  of,  266 
Sugar  of  lead,  340 

see  Lead  acetate 

Sulphates,  solubility  of,  187,  188 
Sulphides,  solubility  of,  178,  180 

test  for,  1 80 
Sulphites,  solubility  of,  183 

test  for,  183 
Sulphur,  176,  177 

allotropic  forms  of,  177,  178 

amorphous,  178 

chemical  relations  of,  178-190 

compared  with  oxygen,  178 

family,  relationships  of  members  of, 
191 

flowers  of,  177 

liquid  forms  of,  177 

monoclinic,  177,  178 

native,  178 

oxygen  derivatives  of,  181 

plastic,  177 

properties  of,  177,  178 

rhombic,  177,  178 

roll,  177,  178 

transition  points  of,  177,  178 
Sulphureted      hydrogen,     see     Hydrogen 

sulphide,  178,  180 
Supercooled  liquids,  177 
Superphosphate  of  lime,  311 
Surface  phenomenon,  29,  66,  68,  69 
Sylvite,  see  Potassium  chloride 
Symbols,  50 
Sympathetic  ink,  364 
Synthesis,  29 


Tables,  list  of, 
Tantalum,  343 
Tartar-emetic,  222,  260 
Tartrates,  solubility  of,  260 
Tellurium,  190 
Temperature,  critical,  199 

critical  solution,  95 

kindling,  56 

scale  according  to  color,  404 
Test  paper,  preparation  of,  59,  106 
Thallium,  328,  332 

properties  and  compounds  of,  332 
Theory,  definition  of,  41 

electronic,  132 

of  ions,  130 


Thermit,  329 
Thermochemistry,  38 

laws  of,  39 

principles  of,  39 

Thermometer,  corrections  for,  383 
Thermometric    readings,    conversion    of, 

383 

Thiocyanates,  237 
Thomas-Gilchrist  process,  356 
Thomson,  J.  J.,  132 
Thorium,  333 
Tin  (stannum),  333,  334,  335,  336 

analytical  reactions  of  salts  of,  338 

properties  and  compounds,  337-339 

table  of  properties  of,  337 
Tincal  (crude  borax),  272 
Tin  cry,  335 
Tincture,  253 
Tin-plate,  336 
Tin-stone  (cassiterite),  335 
Titanium,  333 
Titration,  114,  115 
Triads  of  Dobereiner,  168 
Tridymite,  241 
Triolein,  262,  264 
Tripalmitin,  262,  264 
Tristearin,  262,  264 
Tungsten,  349 

-steel,  350 

Turnbull's  blue,  360 
Turpentine,  106 
Type-metal,  336 

composition  of,  403 

U 

Ultramarine,  330 
Units,  chemical,  52 

of  electrochemistry,  134 
Univalent,  41 

Unsaturated  compounds,  41 
Uraninite,  350 
Uranium,  350 
Urea,  249 


Valence  (quanti  valence),     definitions    of, 

5J>  I3I 

identical  with  electrical  charge  on  the 
ion,  131,  345 

multiple,  51 

table,  174 
Vanadates,  343 
Vanadium,  472 
Van't  Hoff,  quoted,  128,  129 
Vapor  pressure,  62,  84 
Vapor  tension,  lowering  of  the,  97 

of  liquids,  84 

of  water,  tables  of,  389,  390 


INDEX. 


433 


Vaporization  of  water,  heat  of,  79 

Vapor  pressure  of  mercury,  391 

Vaseline,  251 

Verdigris,  294 

Vermillion,  223 

Vinegar,  see  Acid,  acetic,  248,  252 

Vitriol,  blue,  294,  see  Copper  sulphate 

green,  359 

oil  of,  see  Acid,  sulphuric,  184 

white,  319,  359 
Volume,  atomic,  174,  175 

critical,  49 

molar,  52 

W 

Waage    and    Guldberg's    law    of    mass 

action,  120 

Walker,  James,  on  osmotic  pressure,  88 
Washing  soda,  291 
Water,  75 

as  a  solvent,  81 

boiling  temperature  of,  392 

catalytic  properties  of,  105 

chemical  properties  of,  79 

composition  of,  75 

critical  temperature  of,  79 

density  of,  table,  384 

dielectric  constant  of,  132 

dissociation  of,  79 

electrolytic  dissociation  of,  79 

electrolysis  of,  27 

formula  of,  variation  of,  80 

gas,  245 

"hardness"  of,  312 

hydrolytic  action  of,  see  Hydrolysis, 
77,  80,  144 

of  hydra tion  (crystallization),  78 

per  cent,  ionized,  131,  145 

physical  properties  of,  79 

polymerized,  80 

purification  of,  77 

vapor  tension  of,  tables,  389,  390 

volume  of  one  gram  of,  table,  385 
Waters,  chalybeate,  355 
Water  bottle,  construction  of,  376 
Water-glass,  242 


"Weak"  acids  and  bases,  explanation  of 
141-143 

neutralization  of,  143 
Weighings    in    air    reduced    to    vacuum, 

table,  387 
Weight,  atomic,  50 

molecular,  50 
Weights,  atomic,  412-414 

chemically  equivalent,  48 

combining,  48 

equivalent,  48 

molecular,  see  under  each  metal 
Welsbach  mantles,  333 
White  lead,  341 

chemical   principles  involved   in   the 

manufacture  of,  341 
Wines,  bouquet  of,  258 
Witherite,  315 
Wohler,  249 
Wolframite,  349 
Wood,  269 

Wood's  metal,  220,  277 
Wrought  iron,  356 
Wulfenite,  349 


X 


Xenon,  197 


Yeast,  252 

Yellow  prussiate  of  potash,  237 
Ytterbium,  326,  414 
Yttrium,  326,  414 


Zinc,  316,  317,  318 

analytical  reactions  of  salts  of,  320 
blende,  316 

hydrogen  equivalent  of,  71 
properties  and  compounds  of,    318- 

321 

table  of  properties,  317,  318 
vapor  density  determinations  of,  317 

Zincates,  317 

Zircon,  333,  414 

Zirconium,  333,  414 


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